COLLEGE  of  MINING 

DEPARTMENTAL 


BEQUEST  OF 


SAMUELBENEDICTCHRISTY 

PROFESSOR  OF 

MINING  AND   METALLURGY 
1885  -1914 


INTRODUCTION  TO 
METALLURGICAL  CHEMISTRY 


INTRODUCTION   TO 

METALLURGICAL  CHEMISTRY 

FOE  TECHNICAL  STUDENTS 


BY 
J.  H.  STANSBIE,  B.Sc.  (LOND.),  F.I.C. 

ASSOCIATE   OF  MASON  UNIVERSITY   COLLEGE 
AND   LECTURER   IN  THE   BIRMINGHAM   MUNICIPAL  TECHNICAL   SCHOOL 


SECOND  EDITION 


NEW   YOEK 

LONGMANS,     GKEEN     &     CO. 
LONDON  :    EDWARD  ARNOLD 


PREFACE 

THIS  little  book  is  intended  for  the  use  of  technical  students 
who  are  desirous  of  making  a  study  of  the  metals  employed 
for  industrial  purposes.  It  must  be  regarded  strictly  as  a 
preparatory  course,  before  the  study  of  "  metallurgy  "  proper. 
Several  years'  experience  in  teaching  large  classes  in  metal- 
lurgy has  made  it  evident  that  many  who  are  engaged  in 
business  during  the  day,  and  have,  therefore,  but  little  time 
for  study,  are  unwilling  to  devote  several  years  to  a  syste- 
matic course  of  chemistry  before  commencing  the  subject  in 
which  they  are  essentially  interested.  It  has  been  found  that 
two,  or  in  special  cases  three,  evenings  a  week  is  the  usual 
limit  to  the  attendance  of  such  students.  This  permits  of  only 
one  subject  being  taken  at  a  time. 

The  difficulty  has  been  recognised  for  some  years  in  this 
school,  and  at  first  attempts  were  made  to  overcome  it  by 
including  chemical  principles  in  the  ordinary  metallurgy 
classes ;  but  the  short  time  at  the  disposal  of  the  teachers 
made  the  information  fragmentary,  and  of  limited  value.  It 
was  then  decided  to  establish  a  preparatory  class  in  connection 
with  the  metallurgical  department,  in  which  chemical  principles 
could  be  treated  with  special  reference  to  the  wwk  of  that  depart- 
ment ;  and  by  careful  correlation  of  this  with  the  other  classes, 
it  has  been  found  that  the  students  feel  at  once  that  the  work 


vi  PREFACE 

they  are  doing  is  directly  applicable  to  their  particular  branches 
of  industry. 

It  is  becoming  more  evident  every  year  that  the  aim  in  an 
evening  technical  school  should  be  rather  to  increase  the  general 
capability  of  a  large  number  of  students  than  to  turn  out  a  few 
highly-trained  men.  This  end  will  be  best  attained  by  making 
the  subjects  taught  as  self-contained  as  possible. 

The  plan  of  this  book  is  simple.  It  assumes  that  those  who 
use  it  are  practically  interested  in  the  common  metals,  and  that 
they  have  no  further  knowledge  of  their  properties  than  has 
been  obtained  by  ordinary  observation  in  a  workshop  or 
foundry,  so  that  it  is  necessary  to  commence  their  scientific 
treatment  at  the  very  beginning.  The  physical  properties  of 
the  metals  and  their  alloys  being  fully  treated  in  standard 
works  on  metallurgy,  more  than  a  passing  notice  is  not  devoted 
to  this  very  important  part  of  the  study  of  metals.  Brief 
references  to  the  more  useful  properties  will  be  found  in  different 
parts  of  the  text.  The  greater  portions  of  the  chapters  are 
taken  up  with  a  description  of  the  chemical  properties  of  the 
common  metals,  and  in  the  development  of  chemical  principles 
the  metals  are  made  to  take  as  prominent  a  part  as  possible. 
The  non-metals  are,  for  the  purpose  of  this  book,  regarded 
more  as  the  servants  of  the  metals  than  as  their  masters ;  but 
this  is  not  from  any  want  of  appreciation  of  their  importance, 
and  it  is  hoped  that  the  common  ones  are  sufficiently 
described  to  make  them  quite  familiar.  Such  of  the  non- 
metals  as  are  not  of  importance  in  dealing  with  the  chemistry 
of  the  metals  are  either  entirely  omitted  or  only  briefly 
noticed,  further  study  being  left  for  a  more  advanced  stage  of 
the  subject. 

The  practical  work  begins  with  the  study  of  the  effects  of  air 
and  water  on  metals.  This  is  followed  by  a  discussion  of  the 


PEE  FACE  vii 

common  properties  of  matter.  In  dealing  with  the  relation 
between  metals  and  sulphur  the  opportunity  is  taken  of  intro- 
ducing the  preparation  and  properties  of  sulphuric  acid.  This 
is  followed  by  a  description  of  the  nature  and  occurrence  of 
common  salt  and  saltpetre,  from  which,  with  the  aid  of  sul- 
phuric acid,  the  other  common  acids  are  prepared.  By  the 
aid  of  these  compounds  common  substances  of  importance  in 
the  treatment  of  metals  are  prepared  and  described.  Special 
attention  is  paid  to  the  reactions  between  metals  and  acids ; 
and  the  chemical  equivalents  of  some  metals  and  non-metals, 
together  with  their  atomic  weights,  are  considered  at  some 
length.  The  great  importance  of  carbon  in  the  treatment  of 
metals  is  recognised,  and  its  properties  carefully  described. 
This  chapter  is  followed  by  others  in  which  the  important 
subjects  of  combustion  and  reduction  are  somewhat  fully 
considered.  The  pronounced  part  played  by  silicon  and  its 
compounds  in  metallurgical  work  is  also  specially  recog- 
nised. 

Even  a  superficial  examination  of  the  text  will  show  that  the 
book  is  mainly  practical  in  character,  and  that  the  theoretical 
principles  of  the  subject  are  developed  as  far  as  possible  from 
actual  experiments,  to  be  made  either  by  the  student  in  the 
laboratory,  or  by  the  teacher  as  class  demonstrations.  Every 
experiment  has  been  carefully  verified,  and  will  give  good 
results,  if  the  instructions  are  faithfully  followed. 

It  may  appear  to  the  systematic  chemist  that  some  portions 
of  the  subject  have  been  brought  into  undue  prominence,  and 
others  either  scantily  treated  or  omitted  entirely.  This  is  due 
to  the  fact  that  the  book  does  not  profess  to  be  a  work  on 
pure  chemistry,  but  only  an  introduction  to  the  chemistry 
of  the  metals.  It  is  hoped,  however,  that  nothing  of  importance 
has  been  left  out  in  the  development  of  first  principles ; 


viii  PREFACE 

and  that  the  book  will  be  found  useful  both  to  teachers  and 
students. 

The  author's  thanks  are  due  to  Mr.  A.  H.  Hiorns  for  read- 
ing the  proofs,  and  to  Mr.  C.  R.  Clark  for  executing  the 
drawings. 

METALLURGICAL  LABORATORY, 

BIRMINGHAM  MUNICIPAL  TECHNICAL  SCHOOL. 

November,  1903. 


CONTENTS 

CHAPTER  I 
INTRODUCTORY 

PAGES 

Physical  properties  of  bodies— Physical  states— Chemical  change  — 

Energy  1-6 

CHAPTER  II 
METALS   AND   OXYGEN 

The  atmosphere— Oxidation  of  common  metals  when  heated  in  air 
— Burning  of  phosphorus  and  charcoal — Rusting  of  iron — 
Rough  determination  of  the  composition  of  air  -  7-20 

CHAPTER  III 
METALS   AND   WATER 

Oxidation  of  common  metals  by  liquid  and  gaseous  water— Decom- 
position of  water  by  the  electric  current  and  by  heat — Composi- 
tion of  water  by  volume  -  -  21-32 

CHAPTER  IV 
PROPERTIES   OF   MATTER 

Definition  of  matter— Subdivision  of  matter— Composition  of  red 
oxide  of  mercury — Elements  and  compounds — Atoms  and  mole- 
cules— Chemical  symbols — Law  of  Avogadro — Metals  and  non- 
metals— Classification  of  common  elements— Chemical  formulae 
and  equations  -  -  33-54 

ix 


x  CONTENTS 

CHAPTER   V 
METALS  AND   SULPHUR 

PAOES 

Sulphur — Metals  and  sulphur — Formation  of  sulphides — Action  of 
heat  and  air  on  sulphides  —  Preparation  and  properties  of 
sulphuric  acid — Determination  of  formulae  of  compounds  -  55-68 


CHAPTER  VI 
COMMON   ELEMENTS   AND   COMPOUNDS 

Common  salt — Hydrochloric  acid — Chlorine,  bromine,  and  iodine 
— Potassium  chlorate  —  Saltpetre  —  Nitric  acid  —  Ammonium 
chloride  and  ammonia — Caustic  soda  and  potash  -  69-84 


CHAPTER  VII 
METALS  AND  ACIDS 

Reactions  of  common  metals  and  acids — Volume  of  gas  libera  d 
— Metallic  compounds  formed — Hydrogen — Sulphur  dioxide  - 
Oxides  of  nitrogen — Iron  and  acids — Alloys  and  acids  — GenciU- 
tion  of  an  electric  current  by  dissolution  of  zinc  in  dilute 
sulphuric  acid — Decomposition  of  metallic  salts  by  an  electric 
current  -  -  -  -  -  -  85-115 


CHAPTER  VIII 

CHEMICAL  EQUIVALENTS  AND   ATOMIC   WEIGHTS   OF 
COMMON   METALS   AND   NON-METALS 

Determination  of  equivalents  of  common  metals  and  non-metals — 
Vapour  density — Atomic  heat — Isomorphism — Atomic  weights 
— Valency  of  elements  -  116-134 

CHAPTER  IX 
OXIDES,    ACIDS,    AND   SALTS 

Classification  of  oxides  —  Formation  of  salts  —  Crystallization  of 
salts— Nomenclature  of  salts— Chlorides,  nitrates,  sulphates, 
carbonates,  and  other  common  salts — Basicity  of  acids  -  135-162 


CONTENTS 

CHAPTER  X 
CARBON   AND   ITS   COMPOUNDS 

Occurrence  of  carbon— Burning  of  carbon— Composition  of  carbon 
dioxide— Native  and  prepared  carbonates— Action  of  heat  on 
carbonates— Acids  and  carbonates— Carbonic  oxide  :  its  forma- 
tion and  use— Coal— Combustion  of  coal  in  open  and  closed 
grates — "Air  "gas — Coal  gas  -  163-181 

CHAPTER   XI 
REDUCTION 

Action  of  heat  and  of  heat  and  reducing  agents  on  metallic  oxides 
and  other  compounds  —  Carbon,  hydrogen,  carbonic  oxide, 
hydrocarbons,  metals,  and  potassium  cyanide  as  reducing 
agents  —  Reduction  by  reaction  — The  electric  current  as  a 
reducing  agent  -  -  -  182-199 


CHAPTER    XII 
COMBUSTION 

Heat  developed  by  the  burning  of  common  combustible  bodies — 
Ignition-point— Zone  of  combustion— Flame— Luminosity  of 
flame  —  Transformation  of  chemical  energy  into  heat  during 
combustion — Measurement  of  heat  -  200-220 


CHAPTER  XIII 
PHOSPHORUS   AND   ITS   COMPOUNDS 

Common  properties  of  phosphorus — Extraction  of  phosphorus — 
Phosphoric  oxide,  phosphoric  acid,  and  phosphates — Metals 
and  phosphorus  -  -  221-225 

CHAPTER  XIV 

SILICON   AND   ITS   COMPOUNDS   WITH   OXYGEN   AND 
METALS 

Silicon  —  Silica  —  Formation  of  silicates  —  Natural  and  prepared 
silicates — Fusibility  of  silicates — Nomenclature  of  silicates — 
Fireclay— Glass— Effects  of  complexity  on  the  fusibility  of 
silicates— General  classification  -  -  -  226-235 


xii  CONTENTS 

CHAPTER  XV 
WEIGHTS,    MEASURES,    AND   APPARATUS 

PAGES 

The  Metric  System  of  weights  and  measures— Measuring  apparatus 

—The  balance  and  its  use    -  -    236-243 

APPENDIX 

Table  of  elements  and  atomic  weights — Principles  used  in  the  cor- 
rection of  the  volume  of  a  gas  for  variations  in  temperature, 
pressure,  and  amount  of  water  vapour  -  244-247 

INDEX 249-252 


INTRODUCTION  TO 

METALLURGICAL  CHEMISTRY 

CHAPTER  I 
INTRODUCTORY 

THINGS  which  appeal  to  us  from  outside  do  so  through  the 
agency  of  our  senses.  We  are  thus  able  to  recognise  the 
different  objects  of  the  external  world  which  come  under  our 
notice,  and  to  distinguish  between  them  by  means  of  those  of 
their  properties  which  affect  our  sensory  organs. 

Natural  objects  are  called  substances  when  they  are  con- 
sidered with  regard  to  the  quality  of  the  material  of  which 
they  are  made,  without  thinking  about  its  quantity  or  the 
position  which  the  objects  occupy  in  space.  Thus  we  speak  of 
water,  coal,  and  lead  as  substances ;  but  of  a  glass  of  water,  a 
lump  of  coal,  and  a  piece  of  lead  as  bodies.  In  the  first  case 
the  reference  is  general,  and  the  quality  only  of  the  material 
is  recognised ;  while  in  the  second  case  the  quantity  of  the 
material  and  its  position,  as  well  as  its  quality,  are  clearly  in 
mind. 

The  common  and  obvious  characteristics  of  bodies  by  which 
they  are  usually  recognised  belong  for  the  most  part  to  what 
are  called  their  physical  properties.  The  general  appearance 
of  a  body  and  the  various  changes  which  may  take  place  in 
it,  without  altering  the  composition  of  its  substance,  are  due 
to  its  physical  properties.  The  term  "matter"  is  used  to  denote 

1 


2  METALLUKGICAL  CHEMISTKY 

the  substances  of  bodies  generally,  and  any  change  in  this 
matter  causes  a  change  in  the  properties  of  the  bodies  which 
contain  it. 

Physical  States  Of  Matter.— The  substance  of  bodies 
exists  in  three  more  or  less  well-defined  physical  states — the 
solid,  liquid,  and  gaseous.  A  stone  affords  a  good  illustration 
of  a  solid  body,  water  is  the  typical  liquid,  while  the  atmo- 
sphere is  the  most  important  gaseous  body.  Although 
examples  of  solids,  liquids,  and  gases  are  familiar  to  everyone, 
it  is  necessary  to  consider  their  general  properties  sufficiently 
to  be  able  to  distinguish  clearly  between  them. 

Solids. — A  solid  has  a  definite  external  form,  which  it 
retains  as  long  as  it  is  left  to  itself.  If  subjected  to  pressure 
in  any  direction,  it  resists  that  pressure  without  requiring  its 
sides  to  be  supported.  A  change  in  form  may  take  place 
under  sufficient  pressure,  but  when  the  resistance  to  this  change 
is  equal  to  the  pressure  which  is  producing  it,  the  new  form  is 
retained  or  changes  but  very  slowly,  as  long  as  the  pressure  is 
constant.  A  hard  solid,  such  as  steel,  offers  great  resistance 
to  deformation ;  but  a  soft  one,  such  as  rubber,  changes  its 
form  very  rapidly  under  pressure.  Both  bodies,  however, 
regain  their  original  form  on  removal  of  the  pressure,  if  it  has 
not  exceeded  their  power  of  recovery.  The  common  metals, 
with  the  exception  of  mercury,  are  well-defined  solids. 

Liquids. — A  liquid  must,  on  the  other  hand,  have  lateral 
support  if  it  is  to  retain  any  shape,  for  it  has  a  tendency  to 
flow  to  a  lower  level.  The  sides  of  the  vessel  which  contain 
it  supply  this  support,  and  thus  prevent  it  from  flowing.  This 
tendency  to  find  a  lower  level  causes  a  liquid  to  settle  down  in 
the  containing  vessel,  and  present  a  definite  limiting  surface, 
which  is  practically  flat  and  horizontal,  if  the  body  of  liquid 
is  small.  With  liquids  which  wet  the  sides  of  the  vessel  the 
general  surface  is  raised  a  little  near  the  sides  ;  but  with 
liquids  which  do  not  wet  them  a  slight  depression  is 


INTRODUCTORY  8 

observed.     These  effects  are  noticed  when  water  and  mercury 
are  poured  separately  into  dry  test-tubes. 

Some  apparently  solid  bodies,  such  as  cobbler's  wax,  show 
a  decided  tendency  to  flow  if  left  to  themselves;  and  some 
liquids,  such  as  treacle,  are  so  thick  and  viscous  that  they 
flow  very  sluggishly.  Most  of  the  useful  metals,  though  well- 
defined  solids  if  left  to  themselves,  may  be  forced  to  flow 
under  the  hammer,  between  rolls,  or  in  the  press,  and  so  made 
to  take  up  new  forms,  which,  however,  they  retain  after  the 
operation.  A  large  number  of  solids,  including  the  metals, 
may  be  converted  into  liquids  if  made  hot  enough,  and  this 
without  any  change  in  their  composition. 

Gases.— Gases  possess  the  important  and  distinctive 
property  of  indefinite  expansion.  Thus,  if  a  small  body  of  gas 
be  allowed  to  enter  an  empty  vessel,  it  will  expand  and  entirely 
fill  the  vessel,  irrespective  of  its  size.  The  gas  thins  out  in 
the  process,  and  if  it  expands  through  a  very  large  space, 
becomes  highly  attenuated ;  its  surface  is  limited  only  by  the 
sides  of  the  vessel.  This  is  equivalent  to  saying  that  if  the 
pressure  of  the  gas  were  made  very  small,  its  volume  would 
become  greater  than  that  of  any  vessel  into  which  it  could  be 
put.  A  gas  expands  equally  in  all  directions  if  unrestrained, 
and  so  must  completely  fill  any  vessel  which  contains  it. 

Measurable  volumes  of  gases  fulfil,  more  or  less  accurately, 
well-known  laws  when  they  are  heated  and  compressed. 
Also,  gases  mix  together  or  penetrate  each  other  very  readily, 
which  is  no  doubt  due  to  their  property  of  expansion.  Thus 
a  small  quantity  of  an  odorous  gas  liberated  in  a  room  will 
rapidly  expand  through  the  air  of  the  room,  and  may  soon  be 
detected  in  any  part  of  it.  It  is  impossible  to  keep  a  gas  in 
an  open  vessel,  for  some  of  the  gas  escapes  and  air  enters  to 
take  its  place. 

Physical  Change. — Many  bodies  may  be  made  to  assume 
all  three  physical  states  in  succession,  by  heating  or  cooling 
them.  Thus,  water  cooled  below  its  freezing-point  becomes 

1—2 


4  METALLURGICAL  CHEMISTRY 

ice,  and,  if  heated  continuously  at  its  boiling-point,  is  con- 
verted into  steam.  But  the  material  constitution  of  the 
body  remains  the  same ;  there  is  no  alteration  either  in  the 
composition  or  the  weight  of  its  substance.  All  metals  can 
either  be  melted  or  converted  into  vapour  at  varying  tempera- 
tures, which  in  some  cases  are  comparatively  low,  and  in 
others  extremely  high. 

Chemical  Change. — There  is,  however,  another  kind  of 
change  which  bodies  undergo  when  the  necessary  conditions 
are  observed.  It  is  such  as  to  bring  about  an  alteration  in 
the  material  constitution  of  the  body  or  bodies  in  or  between 
which  it  takes  place.  This  is  called  chemical  change,  and  is  so 
characteristic  as  to  be  easily  recognised.  It  depends  upon 
the  chemical  properties  of  the  interacting  bodies,  and  requires 
them  to  be  in  very  close  contact.  Even  then  it  is  often 
necessary  to  start  the  action  by  the  application  of  heat.  For 
example,  gunpowder  may  be  stored  for  years  without  change ; 
but  if  a  light  is  applied  to  it  the  action  commences,  and  the 
change  takes  place  very  rapidly.  The  interaction  of  the 
charcoal,  sulphur,  and  saltpetre,  of  which  the  powder  is  made, 
results  in  the  formation  of  a  very  large  volume  of  gas,  which 
in  its  expansion  produces  some  of  the  observed  effects. 

Energy. — Tt  is  such  a  common  experience  to  see  bodies  in 
motion,  that  we  are  apt  to  take  the  consequent  changes  in 
their  position  as  a  matter  of  course,  and  scarcely  give  a 
thought  to  the  cause  of  the  motion.  When  a  man  lifts  a 
heavy  weight  he  experiences  a  sense  of  muscular  effort,  to 
which  he  ascribes  the  movement  of  the  weight.  The  cause  of 
the  motion  in  this  case  is  called  muscular  force.  In  the 
same  way  any  change  in  the  position  or  motion  of  a  body  is 
due  to  the  action  of  some  kind  of  force  upon  it.  The  study 
of  the  visible  motions  of  bodies  forms  a  branch  of  mechanics 
in  which  the  term  "force"  has  a  definite  meaning,  which, 
however,  need  not  at  present  be  discussed.  On  the  other 
hand,  the  parts  of  a  body  may  be  in  motion,  and  this  motion 


INTRODUCTORY  5 

may  be  increasing  or  decreasing  without  any  motion  of  the 
body  as  a  whole  being  visible.  Thus,  when  a  metal  bar  is 
heated  it  expands,  and  although  this  expansion  is  invisible  to 
the  unaided  eye  it  is  easily  demonstrated  by  actual  measure- 
ment. Now,  it  is  evident  that  the  whole  body  could  not 
expand  unless  its  smallest  parts  or  particles  either  expand 
themselves  or  move  wider  apart.  If  it  is  assumed  that  the 
particles  move  wider  apart,  the  heat  which  is  imparted  to  the 
body  must  be  the  cause  of  their  increased  motion.  If  then  it 
is  necessary  that  force  should  be  exerted  upon  a  body  of 
sensible  size  in  order  to  set  it  in  motion,  it  must  also  be 
necessary  in  the  case  of  a  very  small  body,  or  particle ;  the 
only  difference  in  the  two  cases  is  in  the  relative  magnitudes 
of  the  forces  required. 

It  follows  from  this  that  heat  imparted  to  the  metal  must 
exert  force  upon  its  particles,  and  so  cause  them  to  become 
wider  apart.  There  is  every  reason  to  believe  that  the  ulti- 
mate particles  of  the  metal  are  already  in  motion,  and  that 
the  general  effect  of  the  additional  heat  is  merely  to  increase 
the  rate  of  motion  and  range  of  these  particles.  The  cause  of 
the  increase  may  be  called  heat  force,  but  a  better  name  for 
it  is  heat  ensrgy.  It  must  be  understood  that  the  motion 
of  the  particles  of  a  solid  body  due  to  heating  it,  is  of  a  very 
restricted  character,  and  the  spaces  through  which  they  move 
very  small. 

When  the  charge  of  gunpowder  in  a  cartridge  is  ignited,  the 
bullet  in  front  of  it  is  caused  to  move  very  rapidly,  and  through 
a  considerable  distance.  Now  the  energy  which  brings  about 
this  motion  must  be  stored  up  in  the  powder,  for  it  is  evident 
that  the  mere  blow  of  the  hammer  on  the  detonator  will  not 
account  for  the  effect  produced.  Thus,  when  gunpowder  is 
bought,  the  effective  portion  of  the  purchase  is  the  energy 
which  is  stored  up  in  the  material  of  the  powder.  Further, 
this  energy  can  be  measured  by  the  amount  of  work  it  will  do 
when  set  free.  So  that  it  is  legitimate  to  speak  about  quantity 
of  energy. 


6  METALLUKGICAL  CHEMISTRY 

As  the  energy  of  the  exploding  gunpowder  is  developed, 
the  matter  of  the  powder  itself  undergoes  a  change  in  its 
properties,  and  new  bodies  are  formed.  This  is  a  chemical 
Change,  and  the  energy  which  has  been  transformed  into 
heat,  mechanical  work,  or  other  form  of  energy,  was  chemical 
energy.  It  appears  to  be  associated  in  some  way  with  the 
bodies  between  which  the  change  takes  place,  but  it  is  not 
evident  as  long  as  they  remain  unchanged.  It  is,  however, 
possible  to  measure  one  form  of  energy  indirectly,  while  it  is 
being  transformed  into  another  form  of  energy. 

In  dealing  with  bodies  in  general,  and  the  processes  going 
on  among  them,  two  things  have  to  be  considered  ;  these  are 
the  matter  of  which  the  bodies  are  composed,  and  the 
energy  invariably  associated  with  it.  It  is  most  probable 
that  the  association  of  the  two  makes  the  bodies  what  they 
are,  and  it  is  certain  that  any  alteration  in  this  relation  must 
cause  a  change  of  some  kind  in  the  bodies  themselves.  If 
during  this  change  no  new  bodies  are  formed,  it  is  physical  in 
character,  but  if  new  bodies  are  formed  from  the  old  ones, 
then  it  is  a  chemical  change. 


CHAPTER  II 
METALS  AND  OXYGEN 

THE  atmosphere  which  surrounds  the  earth  is  so  essential  to 
life  that  we  become  familiar  with  it  at  an  early  age,  and 
unconsciously  learn  to  recognise  it  by  its  general  properties. 
Although  invisible,  it  appeals  readily  to  the  sense  of  touch, 
especially  in  a  high  wind.  The  sensation  thus  produced  is 
that  caused  by  a  substance  striking  against  the  exposed  parts 
of  the  body,  and  compels  us  to  admit  that  something  material 
is  appealing  to  us  through  our  senses.  If,  then,  it  is  admitted 
that  the  atmosphere,  although  invisible,  is  matter  or  substance, 
an  investigation  of  those  of  its  properties  which  do  not  appeal 
to  us  directly  may  be  commenced.  Such  work,  if  methodi- 
cally carried  out,  is  sure  to  repay  the  time  and  trouble 
expended  upon  it. 

It  will  be  assumed  that  the  student  knows  very  little  about 
the  atmosphere,  except  that  which  he  has  acquired  by  every- 
day experience.  If  the  experiments  described  below,  or  as 
many  of  them  as  time  and  the  apparatus  at  command  will 
permit,  are  carried  out,  a  considerable  amount  of  information 
about  the  properties  of  the  bodies  under  manipulation,  and 
the  changes  they  undergo  in  the  presence  of  air,  will  be 
obtained.  As  skill  in  manipulation  is  only  acquired  by 
practice,  it  is  well  to  bear  in  mind  that  an  operation,  although 
only  partially  successful,  is  sufficiently  so  if  it  makes  the 
point  it  is  intended  to  illustrate  clear. 

In  the  following  experiments  the  atmosphere  plays  as 
important  a  part  as  the  other  bodies,  which  are  either  liquid 

7 


8  METALLURGICAL  CHEMISTRY 

or  solid,  and  are  therefore  under  direct  observation.  The 
chief  aim  is  to  demonstrate  the  properties  of  the  air,  and  the 
behaviour  of  metals  in  contact  with  it. 

ACTION  OF  HEAT  AND  AIR  ON  METALS. 

Lead. — This  metal  passes  readily  from  the  solid  to  the 
liquid  state  when  heated.  It  has  the  characteristic  properties 
of  the  common  metals.  The  sheet  lead  used  by  plumbers  has 
a  dull  appearance,  but  if  the  surface  is  scraped  with  a  knife, 
it  becomes  bright,  and  shows  the  lustre  common  to  metallic 
bodies.  It  is  soft,  being  readily  cut  with  a  knife,  and  can  be 
beaten  or  rolled  into  sheets:  it  possesses  the  property  of 
malleability  in  a  marked  degree.  A  metal  is  said  to  be 
malleable  when  it  can  be  hammered  or  rolled  into  sheets 
without  cracking  at  the  edges. 

EXP.  1. — Test  the  softness  and  malleability  of  a  piece  of  lead  by 
cutting  it  with  a  knife,  and  by  hammering  it. 

The  surface  of  molten  lead  when  exposed  to  the  atmosphere 
undergoes  a  gradual  change,  which  is  now  to  be  investigated. 

EXP.  2. — Cut  two  squares  of  glazed  paper,  about  3  by  3  in.,  put 
one  in  each  scale-pan  of  the  balance  (Chap.  XV.),  and  make  them 
counterpoise.  Roughly  weigh  about  25  grams  *  of  sheet  lead,  add 
to  it  a  scrap  of  tin,t  about  0*25  gram,  and  weigh  the  whole  accu- 
rately. Transfer  the  lead  to  an  iron  pan  about  6  inches  in  diameter 

(an  ordinary  sand-bath  pan 
will  do),  and  place  the  pan  on 
a  circular  Bunsen  burner,  such 
as  that  shown  in  Fig.  1 ;  or  on 
a  tripod-stand  over  an  ordinary 
Bunsen  flarne.  When  the  lead 
has  melted  drop  in  the  tin,  and 

I  stir  well  with  an  iron  scraper 

pIG    1<  made  of  a  piece  of  narrow  hoop 

iron  about  15  inches  long,  and 

bent  as  shown  in  Fig.  1.     Allow  the  metal  to  remain  at  rest  for  a 

*  Grain  weights  may  be  used  instead  of  grams  :  in  that  case  multiply 
the  number  of  grams  by  15'5  to  convert  into  grains. 

t  The  presence  of  a  small  quantity  of  tin  as  an  impurity  in  the  lead 
increases  the  rapidity  of  the  action  and  renders  the  experiment  less 
tedious  than  with  pure  lead.  Why  it  does  so  need  not  be  considered  now. 


METALS  AND  OXYGEN 

time,  and  watch  the  rapid  change  in  the  appearance  of  its  surface. 
Resume  the  stirring,  and  keep  the  molten  metal  broken  up  into 
globules,  so  as  to  expose  as  large  a  surface  as  possible  to  the  air. 
In  from  thirty  to  forty  minutes  the  metal  will  be  converted  into  a 
coarse  greenish  powder,  when  the  pan  may  be  removed  and  allowed 
to  cool.  When  cold,  transfer  the  powder  with  the  aid  of  a  small 
brush  to  the  counterpoised  paper  upon  which  the  lead  was  weighed, 
and  weigh  it. 

EXAMPLE. — Weight  of  powder  ...     25*615  grams 

„         metal 25-312      „ 

Increase          O303  gram 

This  operation  is  technically  known  as  dPOSSingf,  and  if  a 
longer  time  is  taken  in  carrying  it  out,  a  larger  increase  in 
weight  is  obtained.  But  the  experiment  may  be  considered 
satisfactory  if  any  actual  increase  in  weight  is  shown,  as  it  is 
somewhat  difficult  to  remove  the  whole  of  the  powder  from 
the  pan.  This  is  especially  the  case  when  some  of  the  dross 
has  softened,  through  getting  too  hot,  and  has  stuck  to  the 
pan. 

EXP.  3 — Transfer  the  dross  obtained  in  the  above  experiment  to 
a  clean  porcelain  mortar ;  add  about  10  c.c.  of  water,  and  grind  up 
the  contents  with  the  pestle ;  stir  well  with  the  finger,  and  pour  off 
the  muddy  yellow  liquid  into  a  porcelain  basin  ;  repeat  the  opera- 
tion until  the  water  runs  away  nearly  clear.  Allow  the  basin  to 
stand  for  a  few  minutes  in  order  that  the  suspended  solid  matter 
may  settle.  Then  pour  off  the  clear  liquid,  and  carefully  dry  the 
residue  over  gauze  or  on  the  sand-bath.  When  dry,  remove  the 
residue  from  the  basin,  examine  it,  and  reserve  it  for  future  use. 

The  light  yellow  powder  thus  obtained  corresponds  very 
nearly  to  the  ordinary  massicot  (or  litharge)  in  the  laboratory 
bottle.  It  is  usually  lighter  in  colour  than  the  commercial 
substance,  but  the  longer  the  dressing  stage  is  continued,  the 
deeper  the  colour  of  the  product.  The  washing  operation  is 
called  levigation,  and  if  the  residue  from  it  is  examined,  it  is 
found  to  consist  of  small  particles  of  unchanged  lead,  which 
are  readily  recognised  by  flattening  some  of  the  larger  ones 
on  an  anvil  and  scraping  them.  It  is  thus  evident  that  at 
any  intermediate  stage  in  the  drossing  operation  the  dross 
consists  of  the  yellow  powder  and  the  unchanged  metal.  As 


10  METALLURGICAL  CHEMISTRY 

the  action  proceeds  the  proportion  of  the  unchanged  lead  gets 
smaller,  until  finally  the  whole  of  the  metal  is  converted  into 
massicot. 

EXP.  4. — Put  half  the  massicot  obtained  from  several  experiments 
on  one  of  the  pans  used  for  the  dressing  operation ;  place  the  pan 
on  a  piece  of  gauze  over  a  moderate  Bunsen  flame  ;  stir  the  powder 
occasionally,  and  continue  the  heating  for  twelve  hours.  The  whole 
operation  need  not  be  continuous,  but  the  heating  may  be  carried 
out  from  time  to  time  as  opportunity  occurs. 

At  the  end  of  each  heating  the  colour  of  the  powder  is 
deeper  than  at  the  beginning.  Finally,  it  is  of  a  puce  colour 
when  hot,  and  bright  red  when  cold.  It  will  bear  comparison 
with  the  red  lead  in  the  laboratory  bottle.  The  operation, 
which  is  technically  known  as  colouring",  should  be  shared 
by  those  students  who  have  contributed  massicot  to  the 
charge.  The  pan  may  be  heated  over  a  spare  burner,  and 
needs  very  little  attention. 

The  manufacture  of  massicot  and  red  lead  on  the  large 
scale  is  carried  on  under  similar  conditions  to  the  above  ;  but 
reverberatory  furnaces  are  used,  and  charges  of  a  ton  or  more 
dealt  with  in  one  operation. 

Copper. — This  metal,  which  has  a  characteristic  red  colour, 
is  so  very  malleable  and  ductile  that  it  is  readily  obtained  in 
the  form  of  thin  sheet  or  fine  wire.  If  the  end  of  a  thin 
copper  wire  is  held  in  the  hottest  part  of  the  Bunsen  flame,  it 
melts.  This  indicates  that  the  melting-point  of  the  metal  is 
only  moderately  high. 

EXP.  5. — Cut  a  strip  of  thin  sheet  copper,  about  1-5  inches  wide 
and  4  inches  long,  roll  it  into  a  loose  coil,  and  weigh  it  accurately. 
Put  the  coil  on  a  clay  roasting  dish,  and  place  it  in  a  moderately 
hot  muffle.  Let  it  remain  there  for  about  half  an  hour  ;  then  take 
out  the  dish  with  the  coil  on  it,  and,  when  cold,  reweigh  the  coil. 
Note  the  weight.  Put  the  coil  on  a  piece  of  smooth  paper  to  prevent 
loss,  and  squeeze  it  with  the  fingers.  The  copper  has  probably  all 
disappeared,  and  a  brittle  dark-red  mass  remains.  Grind  it  to  a 
fine  powder  in  a  mortar,  and  transfer  the  powder  to  the  dish.  Put 
the  dish  back  into  the  muffle,  and  stir  the  powder  from  time  to 
time.  Leave  it  in  the  muffle  as  long  as  convenient.  When  cold, 


METALS  AND  OXYGEN  11 

brush  the  powder  carefully  from  the  dish  to  the  piece  of  counter- 
poised paper,  and  weigh  it.  The  temperature  of  the  muffle  should 
not  be  above  a  moderate  red  heat,  or  the  powder  will  soften  and 
stick  to  the  dish. 

EXAMPLE. — Weight  of  powder  after  2nd  heating    3'720  grams 
coil  „     1st        „  3-502      „ 

„    .  3-000      „ 

Increase  in  weight     ...         ...     0'720  gram 

A  piece  of  the  copper  scale,  if  examined  before  being 
powdered,  will  be  found  to  be  blue-black  on  the  outside  and 
dark  red  inside.  It  is  a  mixture  of  two  distinct  bodies,  one 
black  and  the  other  red,  in  varying  proportions.  The  pro- 
portion of  the  black  body  increases  with  the  time  of  exposure. 
The  powder  obtained  by  crushing  the  scale  is  usually  dark 
red  in  colour,  but  after  further  heating  it  becomes  black.  On 
prolonged  heating  in  contact  with  air  the  weight  of  the 
powder  finally  becomes  constant.  A  muffle  furnace  gives  the 
best  result  but  if  one  is  not  available  the  experiment  may  be 
made  as  described  below. 

EXP.  6. — Bend  a  strip  of  sheet  copper  into  the  form  of  a  narrow 
U  ;  weigh  it,  and  suspend  it  on  a  piece  of  clay-pipe  stem  over  a 
Bunsen  flame  in  such  a  position  as  to  get  the  best  heating  effect 
when  the  flame  is  inside  the  U.  Place  a  sheet  of  paper  under  the 
burner,  and  continue  the  heating  for  half  an  hour.  Kemove  the 
strip,  place  it  on  a  convenient  support,  and  cover  it  at  once  with  a 
dry  jar  or  beaker.  When  the  metal  is  cold,  collect  any  scale  which 
may  have  separated  from  the  strip,  and  reweigh  the  whole. 
EXAMPLE. — Weight  of  strip  after  heating  ...  10*580  grams 

„       before     „       ...     10'542      „ 
Increase          0'038  gram 

As  soon  as  the  strip  begins  to  cool,  the  thin  scale  formed  by 
the  action  of  heat  and  air  upon  the  metal  commences  to  peel 
off.  This  is  caused  by  the  rapid  contraction  of  the  cooling 
metal.  The  change  is  due  to  surface  action,  so  that  the  more 
finely  divided  the  metal  is,  the  more  rapidly  it  takes  place. 
With  the  finely-divided  metal  the  maximum  increase  in 
weight  is  obtained  in  a  comparatively  short  time. 

EXP.  7. — Put  a  gram  of  finely-divided  copper,  either  fine  filings 
or  reduced  copper  (Chap.  XI.),  into  a  weighed  porcelain  crucible, 


12  METALLUEGICAL  CHEMISTBY 

and  reweigh  it.     Support  the  crucible  on  a  pipeclay  triangle  over  a 
good  Bunsen  flame.     Observe  the  change  which  takes  place,  and 
continue  the  heating  for  fifteen  minutes.    When  the  crucible  is  cold, 
reweigh  it.     Grind  the  residue  in  a  mortar,  and  examine  it. 
EXAMPLE. — 

Weight  of  crucible  and  copper  after  heating       11-900  grams 
„               „                „           before     „      ...     11-647      „ 
Increase          0*253  gram 

The  difference  between  the  black,  readily-powdered  solid 
and  the  metal  from  which  it  is  formed  is  very  marked. 

Tin. — This  metal  is  soft,  white,  and  very  malleable.  It 
melts  easily,  but  the  change  due  to  the  combined  action  of 
heat  and  air  takes  place  slowly  at  temperatures  below  a  red 
heat.  At  a  red  heat,  however,  the  change  is  fairly  rapid,  and 
is  easily  effected  in  a  muffle.  It  is  more  rapid  when  impure 
metal  is  used.  If  exposed  to  too  high  a  temperature,  the 
metal  burns  and  white  fumes  escape. 

EXP.  8.— Weigh  carefully  3  grams  of  tin,  with  which  is  included 
0*05  gram  of  lead.  Place  the  metal  in  a  roasting  dish,  put  the  dish 
into  the  muffle,  and  keep  it  at  a  moderate  red  heat  for  an  hour. 
Stir  the  metal  occasionally  with  an  iron  scraper,  and  if  it  shows 
signs  of  burning,  draw  the  dish  forward  into  a  cooler  part  of  the 
muffle.  Remove  the  dish,  and,  when  cold,  remove  and  weigh  the 
residue.  Grind  up  the  white  powder  in  a  mortar,  put  it  back  into 
the  dish,  and  heat  it  again  for  half  an  hour,  with  occasional  stirring. 
Weigh  again,  and  note  if  any  increase  is  obtained. 

EXAMPLE. — Weight  of  metal  after  heating  ...     3-793  grams 
before    „        ...     3-000      „ 

Increase 0-793  gram 

Compare  the  white  powder  with  the  metal,  and  with  the  "putty" 
powder  in  the  laboratory  bottle. 

Iron. — This  metal  undergoes  a  somewhat  rapid  change 
when  heated  in  the  air,  and  the  black  scale  that  collects  round 
an  anvil  on  which  red-hot  iron  is  being  forged  is  familiar 
evidence  of  this  change. 

EXP.  9.— Clean  a  few  feet  of  thin  iron  wire  by  rubbing  it  with 
emery-cloth,  and  hammer  or  roll  it  into  a  thin  ribbon.  Weigh 
2  grams  of  the  ribbon,  and  heat  it  to  a  bright  red  heat  on  a  clay 
dish  in  the  muffle  for  half  an  hour.  Allow  the  dish  to  cool,  grind 
up  the  product  in  a  mortar,  and,  if  it  is  not  all  reduced  to  powder, 


METALS  AND  OXYGEN  13 

put  the  residue  back  into  the  muffle,  and  continue  the  heating. 
When  the  whole  has  been  reduced  to  powder,  weigh  it  and  note 
the  increase  in  weight. 

EXAMPLE.— 2  grams  of  iron  increased  to  2*766  grams  of  scale. 

The  general  change  which  many  metals  undergo  when 
heated  in  the  air  may  also  take  place  at  the  ordinary  atmo- 
spheric temperature,  though  much  more  slowly.  It  is  very 
marked  in  the  case  of  iron,  especially  if  the  air  is  moist  and 
contains  acid  vapours.  On  the  other  hand,  tin  will  keep  a 
bright  surface  for  a  long  time  under  ordinary  circumstances, 
and  is  used  as  a  coating  material  for  iron  to  protect  it  from  the 
rusting  action  of  the  atmosphere. 

BURNING  BODIES  AND  Am. 

The  visible  change  which  takes  place  when  a  candle  burns 
is  very  familiar,  for  ordinary  observation  teaches  that  it  dis- 
appear s  during  the  process,  and  that  light  and  heat  are  given 
out      Further  information  is  easily  ob- 
tained by  simple  experiments. 

EXP  10. — Fix  a  short  piece  of  candle  in  a 
deflagrating  spoon,  ignite  it,  and  place  it  in 
a  large  dry  bottle,  as  shown  in  Fig.  2.  At 
first  it  burns  as  brightly  in  the  bottle  as  in 
the  air  outside,  but  the  flame  gradually 
dwindles,  and  finally  disappears.  Also,  the 
sides  of  the  bottle  are  covered  with  moisture. 
Now  pour  a  little  clear  lime-water  into 
another  bottle,  shake  it  up,  and  repeat  the 
burning  of  the  candle  in  this  bottle.  When 
the  flame  has  disappeared,  remove  the  candle 
and  shake  the  bottle  again.  The  lime-water 
is  turned  milky. 

Three  facts  are  demonstrated  by  this  F 

experiment :  (1)  The  burning  of  the 
candle  is  dependent  upon  the  presence  of  air,  for  if  the 
quantity  of  air  is  limited  the  duration  of  the  burning  is 
limited  also;  (2)  moisture  is  formed  during  the  burning; 
(3)  the  air  is  so  changed  by  the  action  that  lime-water  is 
turned  milky  by  it. 


14 


METALLURGICAL  CHEMISTRY 


EXP.  11. — Pour  some  lime-water  into  a  wide-necked  bottle. 
Place  a  piece  of  charcoal  on  the  bowl  of  the  deflagrating  spoon, 
and  hold  the  spoon  in  the  Bunsen  flame  until  the  charcoal  is  red 
hot.  Then  put  it  into  the  bottle,  and  allow  it  to  remain  there  for  a 
minute  or  two.  Remove  the  spoon,  and  shake  the  bottle.  The 
lime-water  is  turned  milky,  thus  indicating  that  the  burning  of 
charcoal  in  air  produces  one  of  the  changes  noticed  during  the 
burning  of  a  candle  in  the  same  body. 

T  he  burning  of  phosphorus  and  the  rusting  of  iron  may  be 
made  to  give  some  important  information  about  the  nature  of 
the  air  in  which  these  actions  take  place. 

Phosphorus. — This  body  is  a  yellow,  waxy-looking  solid, 
which  takes  fire  so  readily  that  it  must  be  kept  in  water. 

EXP.  12. — Fit  up  the  apparatus  shown  in  Fig.  3.  A  is  a  glass  bell 
jar  of  about  2  litres  capacity,  fitted  with  a  well-greased  stopper  or  a 
rubber  bung.  The  jar  is  first  graduated  by  inverting  it  with  the 

stopper  in  position,  and  pour- 
ing in  1,200  c.c.  of  water  from 
a  c.c.  measure  (Chap.  XV.). 
The  water-level  is  marked  on 
the  side  of  the  jar,  300  c.c. 
more  water  added,  and  the 
second  level  also  marked.  The 
jar  is  then  placed  over  water 
in  the  glass  trough  B,  and  so 
supported  that  the  level  of  the 
water  coincides  with  the  1,500 
c.c.  mark  when  the  stopper  is 
out.  In  this  way  1,500  c.c.  of 
air  is  enclosed  in  the  jar  when 
the  stopper  is  in  position,  and 
as  the  water  in  the  jar  forms 
a  movable  bottom  the  air  can 
expand  or  contract  without 

gas  escaping  from  or  entering  the  enclosed  space.  A  small  por- 
celain crucible  containing  a  piece  of  dry  phosphorus  about  the  size 
of  a  pea  is  floated  on  the  water  inside  A.  The  stopper  is  removed, 
the  phosphorus  touched  with  the  hot  end  of  a  piece  of  wire,  and  the 
stopper  rapidly  replaced.  The  phosphorus  burns,  and  a  dense  white 
vapour  fills  the  jar.  In  a  short  time  the  flame  disappears,  and  the 
white  vapour  is  dissolved  by  the  water,  which  gradually  rises  up  the 
jar  to  the  1,200  c.c.  mark.  When  the  white  fumes  have  disappeared, 
it  is  seen  that  about  one-fifth  of  the  original  volume  of  gas  has  been 
removed  by  the  burning  phosphorus,  and  that  the  water  has  risen  in 
the  jar  to  take  its  place, 


B 


FIG.  3. 


METALS  AND  OXYGEN 


15 


Water  is  now  poured  into  B  until  the  level  of  the  liquid  inside  and 
outside  A  is  the  same.  This  will  allow  the  stopper  to  be  removed 
without  any  air  entering  the  jar  through  the  neck,  as  would  be  the 
case  if  the  water-level  inside  had  to  fall.  The  stopper  is  then 
removed,  and  a  burning  candle  lowered  into  the  jar.  The  flame 
disappears,  thus  proving  that  the  residual  gas  is  no  longer  able  to 
support  the  combustion  of  the  candle. 

As  the  phosphorus  burns,  a  body  is  formed  which  is  a 
white  solid  at  the  ordinary  temperature.  This  is  easily  shown 
by  burning  a  piece  of  phosphorus  under  a  dry  bell  jar  standing 
on  a  glass  plate.  As  the  apparatus  cools,  a  snow-white  deposit 
forms  on  the  glass  plate.  With  suitable  apparatus  it  can  be 
proved  that  the  white  solid  is  heavier  than  the  phosphorus 
from  which  it  is  formed.  This  increase  in  weight  comes  from 
the  air,  and  Exp.  12  indicates  clearly  that  as  the  burning 
progresses  the  volume  of  the  air  concerned  in  it  diminishes, 
for  as  the  new  body  is  dissolved  the  water  rises  in  the  jar  to 
take  the  place  of  that  portion  of  the  air  which  took  part  in  its 
formation. 

Iron. — The  rusting  of  iron  takes  place  slowly  under  normal 
circumstances,  but  if  the  metal 
is  finely  divided  and  damp,  the 
action  is  more  rapid. 

EXP.  13.— Fit  up  the  apparatus 
shown  in  Fig.  4.  A  is  a  glass 
cylinder  of  about  500  c.c.  capacity. 
B  is  a  deep  glass  or  earthenware 
dish  in  which  A  can  be  inverted 
and  supported  above  the  bottom 
in  any  convenient  manner.  A  is 
divided  into  five  equal  parts.  To 
do  this,  fill  the  cylinder  with  water, 
pour  it  into  a  measuring  cylinder, 
and  note  the  volume.  Then  pour  the 
water  back  one-fifth  at  a  time,  mark-  J?IG.  4 

ing  the  side  of  the  cylinder  at  the 

level  of  each  fifth.  Cut  a  strip  of  filter-paper  a  little  narrower  than 
the  internal  diameter  of  A ;  thoroughly  wet  the  paper,  and  then  rub 
some  fine  iron  filings  on  both  sides  of  it.  Double  the  prepared  strip, 
and  hang  it  over  a  thin  wire  support  placed  on  C.  Invert  A  over  it, 


!\ 

F 

A 

=-^F. 

SI^~P 

16  METALLURGICAL  CHEMISTRY 

and  pour  water  into  B  until  it  nearly  reaches  the  first  mark  on  A. 
Set  the  whole  aside,  and,  if  possible,  inspect  it  from  time  to  time, 
and  note  the  gradual  rise  of  the  water  in  A.  This  will  go  on  for  a 
day  or  two  and  then  cease  entirely.  If  the  water  rises  one-fifth  up 
the  cylinder,  it  shows  conclusively  that  one-fifth  of  the  original 
volume  of  air  has  been  removed. 

When  the  action  is  finished,  pour  water  into  B  until  the  level  is 
the  same  inside  and  outside  A.  Raise  A  a  little,  and  carefully  draw 
out  the  paper  and  wire -stand,  but  be  careful  to  keep  the  mouth  of 
the  cylinder  below  the  surface  of  the  water  in  B,  or  air  will  enter 
and  spoil  the  experiment.  Pass  a  greased  plate  under  the  mouth  of 
A,  make  it  tight,  lift  out  the  cylinder  with  the  four-fifths  of  gas 
and  one-fifth  of  water  in  it,  and  place  it  mouth  upwards  on  the 
table.  Remove  the  plate,  test  the  gas  with  a  lighted  taper,  and 
quickly  replace  the  plate.  Then  put  the  cylinder  back  into  its 
original  position  in  B,  and  remove  the  plate. 

When  the  paper  used  in  the  above  experiment  is  examined, 
the  particles  of  iron  on  it  show  distinct  signs  of  rusting.  It 
appears,  then,  that  while  the  iron  is  rusting  gaseous  matter  is 
being  absorbed  from  the  air  concerned  in  the  action.  Also,  that 
the  action  ceases  when  about  one-fifth  of  the  total  volume  of 
the  air  has  been  used  up.  This  coincides  with  the  result 
obtained  in  Exp.  12,  and  points  to  the  conclusion  that  only  a 
portion  of  the  air  is  active  in  promoting  the  burning  of  phos- 
phorus and  the  rusting  of  iron. 

EXP.  14. — Weigh  roughly  10  grams  of  red  lead,  put  the  powder 
into  a  piece  of  moderately  thin- walled  combustion- tube,  closed  at 
one  end  and  about  6  inches  long.  Tap  the  tube  until  the  powder 
forms  a  uniform  layer  about  two-thirds  up  the  tube  when  it  is  hori- 
zontal. Fit  the  tube  with  a  cork  and  delivery-tube,  and  support  it 
in  a  horizontal  position,  so  that  the  end  of  the  delivery-tube  can  be 
passed  under  the  mouth  of  A,  Fig.  4.  (See  also  Fig.  14.)  Heat  the 
tube  uniformly  in  the  Bunsen  flame  by  holding  the  burner  in  the 
hand  and  moving  it  about.  Allow  the  air  driven  out  by  the  pre- 
liminary heating  to  escape.  Then  pass  the  end  of  the  delivery-tube 
under  A,  and  strongly  heat  the  red  lead  until  the  jar  is  again  full  of 
gas.  Continue  the  heating,  and  collect  a  little  of  the  gas  in  a  test- 
tube  which  has  been  filled  with  water  and  inverted  over  water  in  B. 

Remove  the  test-tube,  placing  the  thumb  over  its  mouth  to  prevent 
the  escape  of  gas,  raise  the  thumb,  and  insert  a  glowing  splint.  The 
wood  bursts  into  flame,  and  burns  much  more  vigorously  than  in 
the  air. 

Put  the  greased  plate  under  A,  take  out  the  jar,  and  put  it  in  an 


METALS  AND  OXYGEN 


17 


upright  position.  Place  an  empty  jar  by  the  side  of  it,  and  lower  a 
burning  candle,  first  into  one  and  then  into  the  other.  The  candle 
is  found  to  burn  equally  well  in  the  two  jars. 

Clearly,  then,  the  rusting  iron  absorbs  that  part  of  the  air 
which  gives  to  it  the  property  of  supporting  the  burning  of  a 
candle,  and  the  red  lead,  when  heated,  gives  up  a  gas  which 
restores  this  property  to  the  residual  air.  An  examination  of 
the  tube  shows  that  the  red  lead  has  changed  back  to  the 
yellow  substance  from  which  it  was  formed.  It  is  evident  that 
the  red  lead  is  formed  from  massicot  by  the  absorption  of  gas 
from  the  air. 

Mercury. — There  is  another  experiment  which  brings  out 
this  absorption  of  a  definite  gas  from 
the  air  by  the  changing  metal.  But 
it  can  only  be  carried  out  under  ex- 
ceptional circumstances  by  the  ordinary 
student,  as  it  requires  a  moderately 
high  and  uniform  temperature  to  be 
maintained  for  many  hours.  The  fol- 
lowing description  will,  however,  be 
readily  followed,  and  may  be  useful 
for  reference  later. 


EXP.  15. — Fit  a  12-ounce  glass  flask 
with  a  good  cork,  through  which  is  passed 
a  piece  of  glass  tube  about  3  feet  long  and 
{  inch  bore.  Place  it  securely  in  a  deep 
sand-bath  arranged  on  a  retort-stand,  as 
shown  in  Fig.  5.  Eemove  the  cork,  and 
pour  in  mercury  until  the  bottom  of  the 
flask  is  well  covered.  Replace  the  cork 
and  tube,  and  put  a  Bunsen  burner  under 
the  sand-bath.  Heat  the  bath  until  the 
metal  begins  to  boil;  lower  the  flame 
until  the  boiling  ceases,  and  then  raise 
and  lower  it  until  a  little  change,  either 

one  way  or  the  other,  makes  the  mercury  boil  or  takes  it  off 
the  boil.  Keep  the  flask  at  this  temperature  (about  350°  C.)  for 
two  or  three  days.  A  powder  collects  slowly  on  the  surface  of  the 
mercury,  and  when  removed  is  found  to  consist  of  red  scales. 


FIG.  5. 


18  METALLUEGICAL  CHEMISTKY 

If  some  of  this  red  powder  were  used  in  place  of  the  red 
lead  in  Exp.  14,  a  gas  with  exactly  similar  properties  to  that 
collected  from  the  red  lead  would  be  obtained,  and  liquid 
mercury  would  collect  on  the  cool  parts  of  the  tube. 

Mercury  vaporizes  rapidly  at  temperatures  near  its  boiling- 
point,  and  the  principal  use  of  the  long  tube  shown  in  Fig.  5 
is  to  act  as  a  condenser,  and  so  prevent  the  escape  of  metallic 
vapour ;  it  also  serves  for  the  interchange  between  the  air 
inside  and  outside  the  flask,  which  is  necessary  for  the  con- 
tinuance of  the  change. 

The   Energy   of  Chemical   Change.  —  The  changes 

brought  about  during  the  experiments  described  in  this 
chapter  are  sufficiently  characteristic  to  make  it  quite  evident 
that  something  else  happens  besides  an  alteration  in  the 
properties  of  the  reacting  bodies.  We  must  take  notice  of  the 
development  of  heat  which  accompanies  the  disappearance  of 
the  phosphorus  and  carbon  during  their  combustion.  It  may 
be  said  that  the  heat  is  due  to  the  energy  of  combustion,  but 
as  combustion  is  only  a  particular  case  of  chemical  action,  it 
will  be  better  to  substitute  the  more  general  phrase,  "energy  of 
chemical  action."  Clearly,  this  energy  must  be  associated  in 
some  way  with  the  combustible  bodies  and  with  that  part  of 
the  air  which  assists  in  their  burning ;  but  it  is  absent  from 
the  products  of  the  combustion.  It  has  been  transformed  into 
heat,  which  is  simply  another  form  of  energy.  The  energy  of 
chemical  action  is  stored  up  in  the  bodies  which  take  part  in 
the  change,  and,  as  the  action  progresses,  is  converted  largely 
into  heat  in  the  ordinary  processes  of  combustion.  But  other 
forms  of  energy,  such  as  light  and  electricity,  may  also  appear. 
There  is  also  a  very  definite  relation  between  the  quantity  of 
heat  and  any  other  forms  of  energy  developed,  and  the  quantity 
of  chemical  energy  transformed.  There  is  very  little  doubt 
that  no  actual  loss  of  energy  can  possibly  take  place.  As 
one  form  disappears,  other  forms  appear  in  exactly  the  same 
proportion.  This  subject  is  more  fully  dealt  with  in  Chap.  XII. 


METALS  AND  OXYGEN  19 

SUMMARY. 

Several  important  facts  have  been  demonstrated  in  the  fore- 
going experiments. 

The  metals  lead,  copper,  tin,  and  iron,  undergo  very  char- 
acteristic changes  when  heated  in  air,  and  new  bodies  are 
formed,  which  differ  in  such  a  marked  way  from  the  metals 
themselves  as  to  suggest  complete  alteration  in  their  funda- 
mental properties  This  is  to  be  recognised  as  chemical 
change.  It  must  be  carefully  thought  about,  and  every  effort 
made  to  understand  what  it  means,  as  far  as  the  experiments 
and  observations  will  explain  it. 

The  change  in  the  metals  is  accompanied  by  an  increase 
in  weight,  and  it  is  probable  that  in  thinking  about  this  in- 
crease the  mind  will  picture  to  itself  the  addition  to  the  metals 
of  something  which  can  be  weighed. 

The  rusting  of  iron  in  a  confined  volume  of  air  shows  that 
a  definite  part  of  the  air  is  absorbed,  which  is  also  that  part 
which  causes  a  candle  to  burn  in  air.  For  a  candle  will  not 
burn  in  the  residual  gas.  The  restoration  of  this  property  of 
supporting  combustion  to  the  residual  air  by  the  addition  of 
the  gas  obtained  by  heating  red  lead,  seems  to  indicate  that 
the  change  in  the  metals — the  burning  of  a  candle,  charcoal, 
and  phosphorus — are  caused  by  the  same  substance. 

It  is  a  gas,  and  is  called  Oxygen. 

The  increase  in  weight  is  due  to  the  absorption  of  oxygen, 
and  the  action  itself  is  called  oxidation.  When  the  action 
is  very  rapid,  as  in  the  case  of  burning  bodies,  it  is  known  as 
combustion. 

The  residual  gas,  after  removal  of  oxygen  from  air,  will  not 
support  the  combustion  of  ordinary  combustible  bodies.  It  is 
called  Nitrogen. 

However  varied  the  experiments  may  be,  the  volume  of 
oxygen  absorbed  is  always  found  to  be  about  one-fifth  the 
volume  of  the  air  acted  upon.  The  composition  of  pure  air 
is,  then,  roughly  :  Oxygen  i  and  nitrogen  4  by  volume. 

2—2 


20  METALLURGICAL  CHEMISTRY 

Oxygen. — This  is  the  active  constituent  of  the  atmosphere, 
and  as  such  will  claim  considerable  attention  later. 

Nitrogen. — This  gas  is  inactive,  and  for  our  purpose  need 
not  be  considered  further  as  an  isolated  body.  It  will  be 
remembered  as  a  colourless,  transparent,  odourless  gas,  which 
will  not  allow  a  candle  or  other  combustible  body  to  burn  in 
it.  It  reduces  the  activity  of  the  oxygen  of  the  air  by  being 
simply  mixed  with  and  diluting  it.  What  was  formerly 
known  as  pure  atmospheric  nitrogen  is  now  found  to  contain 
a  small  quantity  of  a  heavier  and  even  more  inert  gas  called 
argon. 

The  various  bodies  which  have  been  formed  with  the  co- 
operation of  the  oxygen  of  the  air  in  the  foregoing  experi- 
ments are  called  oxides ;  and,  as  will  be  proved  later,  some 
of  them  are  gases,  which,  as  they  form,  mix  with  the  residual 
air.  If  the  actions  take  place  in  the  open  air,  as  is  often  the 
case,  the  gaseous  oxides  formed  diffuse  into  the  air,  and  form 
part  of  it.  As  such  they  may  be  simply  impurities,  or  they 
may  be  essential  constituents  of  a  good  working  atmosphere. 
This  will  be  considered  in  future  chapters. 

QUESTIONS. 

1.  Descri.be   the    changes  which  take  place  when  lead  01 
copper  is  heated  for  some  time  in  contact  with  air. 

2.  What    do   you   understand   by   the   term    combustible 
body  1     Describe  an  experiment  to  help  your  explanation. 

3.  How  would  you   demonstrate   the   properties  and  pro- 
portions of  the  two  principal  constituents  of  the  air  1 

4.  What  is  meant  by  chemical  change,    arid  how  does  it 
affect  bodies  between  which  it  takes  place  1 

5.  Explain  exactly  what   happens  to  phosphorus   when  it 
burns  in  air. 

6.  How   is   the   process    of   rusting  explained  1     Give   an 
example. 


CHAPTER  III 
METALS  AND  WATEK 

Water. — This  substance  is  the  typical  liquid,  and  is  as  im- 
portant as  it  is  common.  It  comes  next  to  air  as  an  essential 
to  animal  and  vegetable  life.  In  the  pure  state  it  is  a  clear, 
tasteless  liquid,  which  is  quite  colourless  when  viewed  in  bulk. 
It  has  the  property  of  dissolving  a  large  number  of  solids, 
liquids,  and  gases  more  or  less  readily,  and  on  that  account  is 
never  found  perfectly  pure  in  nature.  The  natural  waters 
taken  in  the  order  of  their  purity  are  rain,  river,  spring,  and 
sea  water.  The  impurities  may  be  removed ;  therefore  all 
these  varieties  furnish  samples  of  one  well-defined  invariable 
body — pure  water.  Suspended  impurities,  which  make 
water  look  dirty,  are  easily  removed  by  filtering.  Dissolved 
gases  are  for  the  most  part  got  rid  of  by  boiling  the  water. 
Solids  in  solution  are  left  as  a  residue  on  the  complete 
evaporation  of  the  liquid,  and  the  condensed  steam  is  practi- 
cally pure  water.  Liquid  impurities  are  not  so  readily 
removed,  but  by  repeated  evaporation  their  complete  removal 
is  also  possible.  The  pure  water  used  for  particular  purposes 
is  known  as  distilled  water,  but  ordinary  drinking  water  is 
sufficiently  pure  for  most  purposes.  It  may  be  convenient  to 
use  the  substance  as  ice,  water,  or  steam,  but  whichever  is 
used  the  matter  is  the  same.  Ice  melts  or  water  freezes  at  a 
constant  temperature,  and  water  boils  at  another  temperature, 
which  is  just  as  constant,  if  the  pressure  on  its  surface  is  also 
constant.  These  are  the  standard  "  points  "  for  the  ordinary 
thermometer,  and  on  the  Centigrade  scale,  which  is  commonly 

21 


22 


METALLURGICAL  CHEMISTRY 


used  for  experimental  purposes,  they  are  marked  0°  C.  and 
1 00°  C.  Further  information  about  the  physical  properties  of 
water  will  be  found  in  most  works  on  physics  or  chemistry. 

ACTION  OF  METALS  ON  WATER. 

Sodium. — This  is  a  soft  metal  which  may  be  readily  cut 
with  a  knife.  The  freshly -cut  surface  has  the  ordinary 
metallic  lustre,  and  somewhat  resembles  freshly-scraped  lead 
in  appearance,  but  it  absorbs  oxygen  from  the  air  so  readily 
that  it  tarnishes  very  rapidly.  The  metal  is  usually  kept  in 
some  liquid,  such  as  mineral  naphtha,  which  is  not  acted  upon 
by  it,  so  as  to  exclude  the  air. 

.  EXP.  16. — Fill  a  small  gas  cylinder  of  about  100  c.c.  capacity  with 
water,  and  invert  it  over  water,  as  shown  in  Fig.  6.  Be  sure  that 
no  air  is  left  in  the  cylinder,  or  an  explosion  may  occur.  Cut  a 

piece  of  sodium  about  the  size 
of  a  pea  ;  drop  it  into  the  water 
in  the  trough,  rapidly  place 
over  it  a  gauze  spoon,  and 
force  it  under  the  mouth  of 
the  inverted  cylinder.  Bubbles 
of  gas  escape  through  the 
meshes  of  the  gauze,  rise  in 
the  cylinder,  and  gradually  dis- 
place the  water.  Repeat  with 
similar  pieces  of  the  metal 
until  the  cylinder  is  filled  with 
the  gas.  Sometimes  the  gas 
does  not  get  through  the  gauze 
readily  ;  in  that  case  tilt  the 
spoon  a  little,  and  let  the  metal 
escape,  when  it  will  rise  to  the 
surface  of  the  water  in  the 

cylinder,  and  there  continue  its  action  upon  the  liquid.  Put  a 
greased  plate  over  the  mouth  of  the  cylinder,  and  remove  it  from 
the  water.  Hold  the  cylinder  mouth  downwards,  remove  the  plate, 
and  bring  a  lighted  taper  near.  The  gas  takes  fire,  and  burns 
slowly  with  a  yellowish  flame. 

The  visible  result  of  this  experiment  is  that  the  metal 
dissolves  in  the  water,  and  in  doing  so  liberates  a  combustible 
gas,  which  can  be  collected  and  burnt. 


FIG.  6. 


METALS  AND  WATER 


23 


Magnesium. — This  metal  is  usually  sold  in  the  form  of 
ribbon  or  wire.  It  has  the  ordinary  metallic  lustre,  but 
tarnishes  slowly  in  the  air.  It  is  a  very  light  body  for  a  metal, 
and  when  heated  in  the  Bunsen  flame  unites  with  oxygen  so 
readily  as  to  take  fire  and  burn  with  a  brilliant  light,  leaving 
a  white  residue,  which  contains  magnesium  and  oxygen. 

EXP.  17. — Take  a  piece  of  magnesium  ribbon,  15  inches  long,  and 
make  it  into  a  compact  coil  by  wrapping  it  round  a  pencil.  Draw 
down  a  piece  of  thin  |-inch  combustion-tube  in  the  blowpipe  flame, 
and  bend  the  drawn  part  as  shown  in  Fig.  7.  Push  the  coil  of 
magnesium  into  the  tube  and  draw  off  the  other  end,  bending  it  the 


FIG.  7. 

same  as  before,  and  leaving  the  body  of  the  tube  about  5  inches 
long.  Pass  one  end  through  a  rubber  bung  fitted  in  a  small  flask 
containing  water,  and  attach  a  delivery- tube  to  the  other  by  a  piece 
of  rubber  tubing.  Arrange  the  delivery-tube  so  that  it  can  be 
directed  readily  under  the  mouth  of  a  cylinder  of  about  200  c.c. 
capacity,  which  has  been  filled,  and  inverted  over  water  in  the 
pneumatic  trough.  Boil  the  water  in  the  flask,  and  when  steam  is 
coming  off  freely,  dry  the  body  of  the  combustion-tube  by  carefully 
moving  a  Bunsen  flame  over  it.  When  there  is  no  tendency  for 
water  to  run  back  from  the  upper  bend,  place  the  Bunsen  flame 
directly  under  the  metal  coil,  so  as  to  heat  it  as  strongly  as  possible. 
Be  ready  to  insert  the  open  end  of  the  delivery-tube  under  the 
mouth  of  the  cylinder  directly  the  metal  begins  to  burn.  As  the 
tube  is  hot  from  the  passage  of  steam  through  it,  a  rubber  bung 
slipped  down  to  the  middle  eaves  the  fingers  in  handling  it.  When 


24 


METALLUKGICAL  CHEMISTRY 


the  action  has  ceased  remove  the  delivery-tube,  cover  the  mouth  of 
the  cylinder  with  a  glass  plate,  and  lift  it  out  of  the  trough.  Hold 
it  upside  down,  remove  the  plate,  and  bring  a  lighted  taper  to  the 
mouth.  The  gas  burns  similarly  to  that  obtained  in  the  last  experi- 
ment, but  the  flame  is  not  quite  so  yellow. 

On  examination  the  residue  in  the  combustion-tube  is 
found  to  be  a  white  powder.  It  will  be  useful  to  compare 
this  powder  with  that  obtained  by  igniting  a  piece  of  mag- 
nesium ribbon  in  the  Bunsen  flame,  and  allowing  it  to  burn  in 
the  air. 

[The  apparatus  shown  in  Fig.  7  is  all  that  is  necessary,  but 
it  will  be  safer  to  fix  the  combustion-tube  in  a  clip  to  keep  it 
steady.] 

Iron. — The  general  appearance  and  common  forms  of  this 
metal  are  well  known,  and  it  is  only  necessary  to  point  out 
that  the  sample  used  should  be  clean,  and  show  the  usual 
metallic  lustre. 

EXP.  18. — Clean  a  strip  of  thin  sheet  iron  by  rubbing  it  with 
emery-cloth,  and  cut  it  up  into  narrow  strips  about  1  inch  by  yff  inch. 


Push  a  plug  of  asbestos  to  a  convenient  position  in  a  thin  iron  tube 
about  15  inches  long  and  £  inch  diameter,  and  put  the  clean  iron 
through  the  other  end  of  the  tube  in  sufficient  quantity  to  fill  about 
6  inches  of  it.  Fit  the  ends  of  the  tube  with  rubber  bungs  and 
delivery-tubes,  and  arrange  it  in  a  small  table  furnace,  as  shown  in 


METALS  AND  WATER  25 

Fig.  8.  Wrap  pads  of  blotting-paper  round  the  bungs,  and  keep 
them  saturated  with  water.  This  prevents  the  rubber  from  becoming 
overheated.  Arrange  a  small  flask  for  boiling  water,  so  that  the 
steam  may  be  passed  directly  into  the  tube.  Prepare  to  collect  any 
gas  which  may  be  given  off  from  the  other  delivery-tube.  Heat  the 
iron  tube,  and  boil  the  water.  When  the  steam  comes  into  contact 
with  the  redhot  iron,  gas  is  liberated,  and  collects  in  the  gas  cylinder 
placed  to  receive  it. 

Fig.  8  shows  the  arrangement  of  the  apparatus,  and  when 
about  one-third  of  the  tube  is  filled  with  iron  at  a  good  red 
heat,  the  gas  is  liberated  rapidly,  so  that  several  jars  can  be 
easily  collected.  Closely  adhering  scale  forms  on  the  surface 
of  the  iron.  The  material  of  the  tube  also  helps  the  reaction. 
The  most  suitable  tube  for  the  experiment  is  the  thin,  solid- 
drawn  material  used  for  bicycle  work.  With  this,  if  a  furnace 
is  not  available,  the  reaction  can  be  brought  about  by  heating 
the  tube  in  a  good  Bunsen  flame,  but  in  this  case  the  gas  is 
liberated  slowly,  and  it  would  be  tedious  to  collect  any  con- 
siderable quantity.  By  using  the  foot  bellows  and  blowpipe 
a  better  result  is  obtained. 

Experiments  with  the  Gas.— The  first  jar  of  gas  may 

be  rejected,  as  it  contains  the  air  which  was  present  in  the 
apparatus  at  the  commencement. 

1.  Hold  a  jar  of  the  gas  upside  down ;  remove  the  plate, 
and  bring  a  lighted  taper  near  to  the  mouth.     The  gas  takes 
fire,  and  burns  quietly  with  a  pale-blue  flame,  and  when  the 
burning  taper  is  pushed  through  the  flame  into  the  gas  above, 
it   is  extinguished,  but  is   rekindled  on  being  pulled  back 
through  the  flame.     The  gas  itself  will  burn  (combustible), 
but  will  not  allow  a  taper  to  burn  in  it  (non  supporter  of 
combustion). 

2.  Place  a  jar  with  the  mouth  upwards,  remove  the  plate, 
and  allow  it  to  stand  for  twenty  seconds.     Test  with  a  burn- 
ing taper.     The  gas  has  escaped,  and  the  jar  is  now  full  of 
air. 

3.  Repeat  the  last  experiment,  but  with  the  mouth  of  the 


26  METALLURGICAL  CHEMISTRY 

jar  held  downwards.  Very  little  gas  is  found  to  have  escaped 
when  a  lighted  taper  is  brought  to  the  mouth. 

4.  Hold  an  empty  jar  with  the  mouth  downwards,  and 
bring  a  jar  of  the  gas  into  the  position  shown  in  Fig.  9. 
Keep  them  in  this  position  for  about  half 
a  minute;  separate  the  jars  and  put 
plates  on  them.  On  examination  the 
lower  jar  is  found  to  be  free  of,  and 
the  upper  one  to  contain,  the  greater  part 
of  the  gas,  which  now  burns  with  a  slight 
explosion  when  a  lighted  taper  is  put  to 
the  mouth  of  the  jar.  The  last  three 
experiments  prove  that  the  gas  has  a  very 
great  tendency  to  escape  in  an  upward 
direction  from  vessels  containing  it.  It 
p  q  ^  is,  therefore,  a  very  light  gas. 

5.  Place  a  jar  of  air  and  a  jar  of  the 

gas  mouth  to  mouth,  and  allow  them  to  stand,  one  on  top  of 
the  other,  with  the  air  cylinder  at  the  bottom,  for  fifteen 
minutes.  See  that  the  ground  tops  of  the  cylinders  are  well 
greased,  so  as  to  prevent  any  escape  of  gas.  Separate  them, 
and  put  plates,  first  on  the  bottom  and  then  on  the  top 
cylinder.  Test  both  with  a  lighted  taper.  Equally  loud 
explosions  in  the  two  jars  show  that  the  mixture  of  the  air 
and  the  gas  is  perfect. 

The  name  of  the  gas  is  Hydrogen.  It  burns  with  a  pale 
blue  flame.  It  will  not  support  the  combustion  of  a  taper,  or 
of  similar  bodies.  It  is  a  very  light  gas,  and  can  be  poured 
upwards.  It  is  colourless,  transparent,  and  odourless  when 
pure  ;  but  as  collected  in  the  ordinary  way  it  has  an  odour, 
due  to  impurities  derived  from  the  materials  used  in  its  pre- 
paration. The  property  of  gases  which  enables  hydrogen  and 
air  to  mix  together  perfectly,  although  the  vessel  containing 
the  lighter  gas  is  placed  at  the  top,  is  known  as  diffusion, 
and  is  a  common  property  of  all  gases.  Any  number  of  gases 
will  diffuse  into  each  other,  and  make  a  uniform  mixture, 


METALS  AND  WATER  27 

providing  they  do  not  act  upon  each  other  chemically.  The 
composition  of  a  mixture  of  gases  is  the  same  throughout. 
The  heavier  gases  do  not  usually  separate  to  the  bottom 
and  the  lighter  to  the  top,  however  long  the  mixture  may  be 
kept. 

Action  of  the  Electric  Current  on  Water.— The  use 

of  the  electric  current  for  lighting  purposes  is  now  more  or 
less  familiar  to  everyone,  and  the  presence  of  the  white-hot 
filament  in  the  incandescent  lamp  is  readily  noticeable.  The 
heat  and  light  developed  in  the  lamp  are  due  to  the  resistance 
offered  by  the  carbon  thread  joining  the  terminals  of  the 
lamp  to  the  current  passing  through  it.  These  terminals  are 
connected  with  copper  wires,  which  are  connected  in  turn 
with  the  electric  mains  through  which  the  principal  current  is 
passing,  and  a  small  quantity  is  thus  taken  off  to  be  used  in 
the  lamp.  It  is  an  easy  matter  to  cut  one  of  the  wires  con- 
nected directly  with  the  lamp,  and  to  connect  the  severed 
ends  with  a  vessel  containing  water,  so  that  if  the  current  is 
to  pass  through  the  lamp  it  must  also  pass  through  the  water. 
Now  it  is  found  that  perfectly  pure  water  prevents  the  passage 
of  the  current,  and  no  visible  effect  upon  the  liquid  is  noticed; 
but  if  a  little  acid  is  added  to  the  water  the  current  passes  and 
the  lamp  glows.  The  nature  of  electricity  and  the  cause  of 
its  flow  through  the  wire  and  the  water  cannot  be  discussed 
at  this  stage ;  but  it  will  be  sufficient  for  our  present  purpose 
to  say  that  it  can  be  drawn  off  and  used  to  bring  about  heat 
effects  and  chemical  change,  just  as  readily  as  coal-gas  can  be 
passed  into  a  Bunsen  burner,  and  the  heat  developed  by  its 
burning  used  for  a  similar  purpose. 

The  apparatus  used  to  contain  the  water  is  called  a  volta- 
meter, and  consists  essentially  of  a  glass  vessel  through  the 
bottom  of  which  two  platinum  wires,  with  narrow  strips  of 
platinum  attached,  are  fused.  These  are  shown  in  Fig.  10. 
The  cut  ends  of  the  lamp  wire  are  joined  to  the  ends  of  the 
platinum  wires  outside  the  vessel,  and  acidulated  water  is 


28 


METALLURGICAL  CHEMISTRY 


poured  into  the  vessel  to  fill  in  the  gap  between  the  two 
platinum  plates.  The  lamp  glows,  and  bubbles  of  gas  collect 
on  the  plates,  rise  to  the  surface  of  the 
\              /  water,  and  escape  into  the  air.     The  gas 
\nn/  thus   liberated  can   be  collected  for  ex- 
it?!   \T  K     ffi|  animation  by  filling  two  tubes  with  water, 
p      -Q  and  inverting  them  over  the  plates. 

EXP.  19. — Fit  up  the  apparatus  shown  in 

Fig.  11.  This  voltameter  is  a  very  convenient  form  of  the  apparatus. 
The  platinum  plates,  P  and  P,  are  in  the  tubes  A  and  B,  and  their 
wires  pass  to  the  outside  through  the  glass.  The  glass  taps  at  the  top 
of  A  and  B  are  opened,  and  water  containing  a  little  acid  is  poured 
down  C  until  it  rises  to  the  level  of  the  taps,  which  are  then  closed. 
The  water  is  now  at  the  same  level  in  the  three  tubes,  and  on 
making  the  connec- 
tions with  the  lamp 
wires  the  current 
passes  through  the 
water,  and  gases  are 
liberated  from  the  plat- 
inum plates.  The  bub- 
bles of  gas  rise  in  A 
and  B,  and  force  the 
water  back  into  C.  It 
is  at  once  noticed  that 
the  volumes  of  the 
gases  collecting  in  the 
tubes  are  not  equal, 
and  when  sufficient 
has  collected  it  is  seen 
that  the  simple  rela- 
tion of  1  to  2  is  main- 
tained as  long  as  the 
action  continues.  When 
the  level  of  the  water 
in  A  nearly  reaches  the 


MAINS 

FIG.  11. 


top    of    the    platinum 
plate,   the    connection 

is  broken,  and  the  liberation  of  gas  ceases.  The  larger  volume  of 
gas  is  now  tested  by  cautiously  opening  the  tap,  and  at  the  same 
time  holding  a  lighted  taper  near.  The  issuing  gas,  driven  out  by 
the  fall  of  the  water  level  in  C,  takes  fire  and  burns  with  a  pale  blue 
flame.  When  the  gas  in  B  is  allowed  to  escape  by  opening  the  tap, 
and  a  glowing  splint  is  held  near  the  jet,  the  wood  bursts  into  flame, 


METALS  AND  WATEK 


29 


and  burns  vigorously.  These  tests  are  sufficient  to  indicate  the 
nature  of  the  gases.  The  larger  volume  of  gas  is  hydrogen,  and 
the  smaller  one  oxygen. 

If  the  taps  are  turned  off  and  the  connections  reversed,  the 
hydrogen  collects  in  B,  and  the  oxygen  in  A.  This  shows 
that  the  position  of  the  gases  in  the  apparatus  depends  upon 
the  direction  of  the  current  through  the  liquid.  It  is  usual  to 
state  that  the  oxygen  is  liberated  on  the  plate  from  which 
the  current  enters  the  liquid,  and  the  hydrogen  on  the  plate 
through  which  it  leaves  the  liquid.  The  acid  put  into  the 
water  at  the  beginning  of  the  experiment  is  there  at  the  end, 
unaltered  either  in  quality  or  quantity.  This  shows  that  the 
gases  must  come  from  the  water.  Thus  we  have  additional 
evidence  that  water  is  a  combination  of  oxygen  and  hydrogen, 
and,  further,  that  they  are  present  in  the  proportion  of 
1  volume  of  oxygen  to  2  volumes  of  hydrogen.  A  simpler 
form  of  the  apparatus  similar  to  that  shown  in  Fig.  10,  in 
which  two  tubes,  filled  with  water,  are  inverted  over  the 
platinum  plates  in  the  open  vessel,  may  be  used.  It  can  be 
worked  readily  with  two  Bunsen  cells  if  the  platinum  plates 
are  not  too  far  apart. 

The  two  gases  can  also  be  collected  together,  and  a  very 
explosive  oxy -hydrogen  mix- 
ture obtained. 

EXP.  20. — A  is  a  small  bottle 
containing  acidified  water,  and 
fitted  with  a  rubber  bung, 
through  which  pass  the  two 
platinum  wires  tipped  with 
platinum  plates,  and  a  delivery- 
tube  leading  under  the  inverted 
cylinder  B.  When  the  current 
is  passed  through  the  water  in 
A,  the  mixed  gases  are  given  off 
and  collect  in  B.  If,  when  a  small  cylinderful  has  been  collected, 
it  is  removed  and  a  lighted  taper  put  to  the  mouth,  a  loud  explosion 
takes  place.  The  gases  combine  together  again,  and  water  is 
reformed. 

If  the  lighting  current  is  not  available  for  these  experi- 


FIG.  12. 


METALLURGICAL  CHEMISTRY 


ments  a  voltaic  battery  of  three  or  four  cells  may  be  used. 
(See  Chap.  VII.) 

Action  Of  Heat  on  Water.— When  steam  is  raised  to  a 
high  temperature,  it  is  split  up  into  its  constituents,  oxygen 
and  hydrogen,  but  as  the  temperature  falls  recombination 
takes  place.  If,  however,  the  gases  as  they  are  liberated  are 
allowed  to  mix  with  some  neutral  gas,  the  recombination  can 
be  prevented,  and  evidence  of  the  decomposition  obtained. 

EXP.  21.— A  small  long-necked  flask,  A,  is  fitted  with  a  bung  and 
delivery-tube.  A  spiral  of  thin  platinum  wire  is  stretched  between 

the  lower  ends  of 
two  stout  copper 
wires  which  pass 
through  the  bung, 
and  terminate  in  the 
middle  of  the  body 
of  the  flask.  The 
water  in  A  is  kept 
gently  boiling,  and 
the  current  is  sent 
through  the  spiral, 
which  is  thus  raised 
to  a  bright  red  heat 
by  the  resistance  of 
the  wire  to  the  pas- 
sage of  the  current. 
FIG.  13.  The  steam  in  direct 

contact      with      the 

spiral  is  decomposed,  and  the  liberated  gases  diffuse  into  the  main 
body  of  the  steam.  Bubbles  of  gas  issue  slowly  from  the  delivery- 
tube,  and  collect  in  B.  On  testing  the  evolved  gas,  it  is  found  to 
be  explosive,  and  to  resemble  that  obtained  in  Exp.  19.  The  intense 
heat  of  the  wire  decomposes  the  water ;  the  current  makes  the  wire 
hot,  but  takes  no  further  part  in  the  action. 


B 


SUMMARY. 

The  experiments  of  this  chapter  all  lead  to  the  conclusion 
that  water  is  made  up  of  two  gases,  oxygen  and  hydrogen. 
In  its  decomposition  by  the  electric  current  the  information 
obtained  is  very  precise,  for  the  water  is  proved  to  contain 
the  two  gases  in  the  proportion  of  2  volumes  of  hydrogen 


METALS  AND  WATEK  31 

to  1  volume  of  oxygen.  Also,  when  the  gases  are  mixed 
together  in  these  proportions,  and  a  lighted  taper  is  put 
into  the  mixture,  recombination  takes  place  with  great 
violence. 

Hydrogen  has  very  characteristic  properties.  It  burns  with 
a  pale  blue  flame,  and,  as  will  be  proved  later,  water  is  formed 
during  the  combustion.  It  is  a  very  light  gas,  being  readily 
poured  upwards  from  one  vessel  to  another.  It  also  diffuses 
into  other  gases  very  readily.  The  heavier  a  gas  is  the  more 
slowly  it  diffuses.  The  general  law  states  that  the  rate  of 
diffusion  of  gases  varies  inversely  as  the  square  roots  of  their 
densities.  Hydrogen  is  the  lightest  body  known,  and  has, 
therefore,  the  highest  rate  of  diffusion. 

Ice  when  heated  melts  to  water;  water  when  boiled  is 
converted  into  steam  ;  steam  when  very  strongly  heated  is 
split  up  into  oxygen  and  hydrogen.  The  first  two  changes 
are  purely  physical,  for  the  actual  composition  of  the  substance 
is  unchanged.  The  last  is  a  change  in  composition,  for  other 
forms  of  matter  appear  in  proportion  as  the  water  disappears. 
This  is  a  chemical  change. 

The  common  metals,  with  a  few  exceptions,  such  as  sodium 
and  potassium,  do  not  attack  pure  water  at  ordinary  atmo- 
spheric temperatures,  and  may  be  left  in  contact  with  it  for  a 
long  time  without  appreciable  change.  But  this  is  not  the 
case  with  natural  water,  which  always  contains  impurities, 
and  these  often  act  either  by  themselves,  or  help  to  bring  about 
an  action  between  the  metal  and  water.  The  dissolved  oxygen 
and  carbonic  acid  gases,  which  are  absorbed  from  the  air  by 
water  when  exposed  to  it,  are  active  agents  in  the  case  of 
several  metals.  Thus  iron  rusts  readily,  and  lead  is  acted 
upon  somewhat  rapidly  by  such  impure  water.  The  presence 
of  solids  in  solution  is  also  effective  in  promoting  change. 
This  is  noticed  in  the  corrosion  of  those  metal  portions  of 
ships  which  are  exposed  to  the  action  of  sea-water.  A  large 
proportion  of  the  common  metals,  when  at  a  red  heat,  attack 
water  vapour  readily,  and  in  all  such  cases  oxides  of  the 


32  METALLUKGICAL  CHEMISTKY 

metals  are  formed.     Copper,  lead,  and  tin  act  very  slightly,  if 
at  all,  even  at  a  bright-red  heat. 

QUESTIONS. 

1.  How    would    you    obtain    a    sample    of    pure    water1? 
Describe  its  chief  physical  properties. 

2.  What  kind  of  body  is  sodium?     State  generally  what 
takes  place  when  it  is  thrown  on  to  water. 

3.  How  may  hydrogen  gas  be  obtained  in  sufficient  quan- 
tity to  test  its  common  properties  ?     Name  these  properties. 

4.  Make   a  sketch  of  a   simple  form  of   voltameter,  and 
explain  how  you  would  use  it. 

5.  Can  an  analogy  be  drawn  between  the  action  of  heat 
and  of  an  electric  current  upon  water  ? 

6.  Explain  in  a  simple  way  the  difference  between  physical 
and  chemical  change. 

7.  Describe  the  property  of  diffusion  of  gases. 

8.  What  is  the  nature  of  the  bodies  formed  when  metals 
act  upon  water  vapour  ? 


CHAPTER  IV 
PEOPEETIES  OF  MATTEE 

MATTER  is  the  general  name  given  to  the  stuff  or  substance  of 
which  all  bodies  are  made  up.  A  little  thought  will  make  it 
clear  that  the  total  number  of  different  bodies  is  very  large, 
and  it  would  seem  an  almost  endless  task  to  attempt  their 
classification.  Very  little  difficulty  is,  however,  experienced 
in  arranging  the  majority  of  bodies  in  three  great  divisions— 
the  animal,  vegetable,  and  mineral  kingdoms. 

Mass  and  Density.— It  may  be  stated  that  the  "mass," 
or  quantity  of  matter  in  a  body,  is  a  property  of  the  matter 
which  accompanies  it  without  loss  or  gain  through  every 
possible  change  which  the  body  can  undergo.  Mass  is  most 
easily  estimated  in  practice  by  weighing,  so  that  the  weight 
of  a  body  may  be  taken  as  the  measure  of  the  amount  of 
matter  in  it ;  but  it  must  be  remembered  that  the  terms 
weight  and  mass  have  somewhat  different  meanings.  The 
weight  of  a  body,  as  measured  on  a  spring-balance,  varies 
according  to  the  position  of  the  body  with  regard  to  the 
earth,  for  the  pull  of  the  earth  upon  the  body  varies  with  the 
distance  of  its  centre  from  the  centre  of  the  earth.  The  force 
of  gravity,  which  is  the  cause  of  the  weight  of  a  body, 
decreases  as  the  distance  of  the  body  from  the  earth's  centre 
increases,  so  that  a  given  body  would  stretch  the  spring  of  the 
balance  more  at  a  place  near  one  of  the  poles  than  it  would  at 
a  place  nearer  the  equator,  as  at  the  latter  the  distance  of  the 
earth's  centre  from  it  would  be  greater.  Thus,  it  may  be  said 

33  3 


34  METALLUKGICAL  CHEMISTEY 

that  the  mass  of  a  body  depends  solely  upon  the  amount  of 
matter  in  it ;  and  its  weight  upon  the  latter,  together  with  the 
proximity  of  other  bodies  to  it.  But  in  ordinary  weighing  on 
a  beam-balance,  the  body  placed  in  one  pan  is  counterpoised 
by  the  weights  placed  in  the  other,  so  that  any  change  in  the 
force  of  gravity  affects  the  body  and  the  weights  equally,  and 
for  practical  purposes  a  sufficiently  true  estimate  of  the  mass 
of  a  body  is  obtained  by  determining  its  weight  in  the  usual 
way. 

The  density  of  a  body  is  usually  stated  in  terms  of  the 
quantity  of  matter  in  a  unit  volume  of  the  body  as  determined 
by  weighing.  But  to  be  accurate,  it  is  necessary  to  state  the 
temperature  of  the  body  in  the  case  of  solids  and  liquids,  and 
both  the  temperature  and  the  pressure  in  the  case  of  gases  or 
vapours  at  the  time  of  weighing.  All  bodies  expand  when 
heated  and  contract  when  cooled ;  and  variations  in  the 
pressure  to  which  a  body  of  gas  is  subjected  cause  changes  in 
its  volume  which  cannot  be  neglected  even  in  rough  experi- 
ments ;  but  changes  in  the  volume  of  solids  and  liquids  from 
the  same  cause  are  so  small  as  to  be  negligible  in  the  ordinary 
way.  Thus  the  density  of  gold  is  19-3  grams  per  cubic  centi- 
metre at  15°  C. ;  that  of  water  1  gram  per  cubic  centimetre  at 
4°  C.  ;  and  that  of  oxygen  OO0143  gram  per  cubic  centimetre 
at  0°  C.  and  76  centimetres  barometric  pressure.  (See 
Chap.  XV.  and  Appendix.) 

The  weighing  of  solids  and  liquids  is  a  comparatively  easy 
operation  ;  but  in  the  case  of  gases  it  is  much  more  difficult 
and  tedious.  In  the  method  of  weighing  gases  adopted  by 
Regnault,  two  equal  glass  globes  fitted  with  stopcocks,  and 
having  a  capacity  of  several  litres,  are  used.  One  of  the 
globes  is  exhausted  by  means  of  a  good  air-pump,  and  then 
suspended  from  a  hook  fixed  to  the  bottom  of  the  weight  pan 
of  a  very  sensitive  balance.  This  globe  serves  as  a  counter- 
poise for  the  other  one,  whi.h  is  also  completely  exhausted 
and  similarly  suspended  from  the  weighing  pan  of  the  balance. 
The  globes  are  now  carefully  counterpoised,  and  the  one  to 


PROPEETIES  OF  MATTEE  35 

be  used  for  the  actual  weighing  is  filled  with  the  pure  gas  at 
a  pressure  slightly  gre  tter  than  that  of  the  atmosphere  out- 
side. It  is  again  hung  from  the  scale  pan,  and  the  stopcock 
op  ned  and  closed  rapidly  to  bring  the  gas  inside  to  the  same 
pressure  as  the  air  outside.  The  temperatut  e  of  the  balance 
case  in  which  the  globes  hang,  and  the  barometric  pressure  of 
the  air,  are  the  a  noted,  and  the  weighing  completed.  If, 
then,  the  capacity  of  the  globe,  which  determines  the  volume 
of  the  gas,  is  known,  and  this  volume  is  reduced  by  calculation 
to  that  which  the  gas  would  occupy  at  the  normal  temperature 
and  pressure  (N.T.P.),  the  weight  of  unit  volume  of  the  gas 
under  normal  conditions  is  easily  calculated.  (See  Appendix.) 
Regnault  found  that : 

1  c.c.  of  hydrogen  weighs  0-0000896  gram 

oxygen  „       0-0014304     „ 

„         air  „       0-0012936     „ 

In  the  determination  of  the  density  of  the  vapours  of  solids 
and  liquids  which  can  be  vaporized  by  heat  at  workable  tem- 
peratures, it  is  not  necessary  to  actually  weigh  the  vapours, 
for  the  solid  or  liquid  may  be  weighed  and  then  converted 
into  vapour,  and  its  volume,  temperature,  and  pressure 
observed.  The  volume  of  the  known  weighs  of  vapour  at 
N.T.P.  is  then  easily  calculated,  and  the  density  of  the  vapour 
obtained.  Sometimes  it  is  preferred  to  weigh  the  liquid  or 
solid  after  it  has  been  vaporized  and  condensed  again.  The 
accuracy  of  these  determinations  depends  upon  the  fact  that  no 
change  in  the  mass  of  the  body  takes  place  during  the  conversion 
into  vapour. 

Conservation  Of  Matter.— All  practical  experience  leads 
to  the  belief  that,  although  bodies  in  general  may  be  made 
to  undergo  a  variety  of  changes,  the  matter  in  them  cannot 
be  destroyed.  It  preserves  its  quantity,  and  passes  without 
loss  through  every  known  process  to  which  bodies  can  be 
subjected.  Nor  can  matter  be  created.  The  changes  that 
take  place  in  bodies  which  are  passing  through  a  given  process 

3—2 


36  METALLURGICAL  CHEMISTRY 

may  be  very  complicated,  but  there  is  no  evidence  of  gain 
in  the  quantity  of  matter  they  contain.  These  statements, 
that  matter  can  be  neither  created  nor  destroyed,  refer 
essentially  to  its  quantity,  and  are  supported  by  every  ex- 
periment in  which  the  weights  of  the  changing  bodies  are 
considered.  Any  apparent  increase  in  the  weight  of  a  body 
which  has  undergone  change  can  be  traced  to  the  transference 
of  matter  to  it  from  some  other  body  that  has  taken  part 
in  the  change,  and  has  suffered  an  equal  decrease.  This  is 
exemplified  in  the  oxidation  of  metals  in  air.  The  general 
principle  is  usually  stated  under  the  title  of  the  conservation 
of  matter. 

Divisibility  of  Matter.  —It  may  be  taken  as  an  admitted 
fact  that  any  mineral  body  can  be  divided  into  parts,  and 
that  these  parts,  under  exceptional  circumstances,  may  be 
exceedingly  small.  All  ordinary  experience  is  in  favour  of 
this  statement,  for  even  the  diamond,  which  is  the  hardest 
known  body,  can  be  ground  to  a  fine  powder  on  the  lapidary's 
wheel ;  but  it  is  diamond  cut  diamond,  for  diamond  dust 
must  be  used  in  the  process.  But  however  finely  divided 
a  body  may  be,  its  ultimate  particles  must  have  a  definite 
size,  although  they  may  be  excessively  small.  Each  particle, 
however  small,  must  have  a  definite  volume  ;  that  is,  the 
mass  of  matter  in  it  must  occupy  a  finite  space  to  the 
exclusion  of  all  other  matter  frcm  that  space.  This  is  some- 
times expressed  by  saying  that  matter  possesses  volume. 
Another  general  experience  of  all  bodies  that  can  be  handled 
is  that  they  possess  weight,  and  that  this  property  is  in 
proportion  to  the  amount  of  matter  in  them.  If,  then,  these 
general  properties  of  matter  are  admitted,  irrespective  of  the 
size  of  the  body,  the  smallest  possible  particles  of  a  body  must 
occupy  a  definite  volume,  and  contain  a  definite  quantity  of 
the  matter  of  the  body. 

Subdivision  of  Gold.  —  This  metal  has  many  useful 
properties,  and  one  of  the  most  impoitant  is  its  great  malle- 


PKOPEKTIES  OF  MATTEK  37 

ability.  It  may  be  beaten  out  so  thin  that  a  sheet  having  an 
area  of  50  square  inches  weighs  only  1  grain.  A  square  inch 
of  this  gold  leaf  is  readily  cemented  to  a  sheet  of  glass,  cut 
into  100  strips  by  a  sharp  tool,  and  each  of  these  divided  into 
100  parts  by  cutting  across  the  square  again  at  right  angles 
to  the  first  direction.  That  is,  -Jg-  grain  of  gold  may  be 
cut  into  100x100  =  10,000  small  pieces,  or  1  grain  into 
500,000  small  pieces  ;  and  these  can  be  seen  with  a  good 
pocket  lens.  But  this  is  not  by  any  means  the  limit  to  the 
subdivision  of  the  gold ;  for  with  a  special  machine  it  is 
possible  to  rule  10,000  lines  across  the  square  inch  of  leaf, 
and,  by  ruling  again  at  right  angles  to  the  first  direction,  to 
cut  the  whole  square  into  10,000  x  10,000  =  100,000,000  small 
pieces.  Thus  a  single  grain  of  gold  can  be  mechanically 
divided  up  into  100,000,000x50  =  5,000,000,000  pieces,  and 
these  pieces  can  be-  seen  under  a  high-power  microscope. 
It  may,  then,  be  said  that  is  possible  to  get  a  piece  of  gold 
which  is  only  ^oooWorr  grain  in  weight,  and  is  still 
visible  to  the  aided  eye.  This  division,  extreme  though 
it  may  appear,  is  not  the  limit  of  possibility,  for  still  finer 
work  can  be  done. 

Dissolution  of  Common  Salt.  —  Compare  the  above 
operation,  which  depends  upon  great  mechanical  skill,  with 
the  simple  process  of  dissolving  common  salt  in  water.  The 
solid  disappears  rapidly,  and  the  most  powerful  microscope 
fails  to  detect  the  particles  into  which  it  is  divided.  That  it 
is  simply  divided  is  readily  proved.  The  liquid  tastes  of  the 
salt,  and  if  the  water  is  evaporated  the  particles  run  together 
again,  and  the  &olid  salt  is  recovered.  Very  little  is  known 
about  the  actual  size  of  the  salt  particles  in  the  solution,  but 
it  is  clear  that  they  are  smaller  than  a  particle  of  gold  weigh- 
ing sinrWoinMny  grain.  Also,  there  is  evidence  in  favour  of 
the  statement  that  when  a  small  quantity  only  of  the  salt 
is  present  in  the  solution,  its  separate  particles  are  as 
small  as  they  can  possibly  be.  As  will  be  explained  later, 


38  METALLURGICAL  CHEMISTRY 

the  term  molecule,  which  means  a  small  mass,  is  used  to 
designate  the  smallest  possible  particle  of  the  salt  which  can 
have  a  free  existence.  The  molecules  of  salt  are  exessively 
small,  but  none  the  less  definite,  and  go  to  make  up  the  mass 
of  a  lump  of  salt  of  sensible  size. 

Gases. — These  bodies  are  for  the  most  part  invisible,  but 
it  is  certain  that  the  body  of  any  gas  is  made  up  of  an  enormous 
number  of  very  small  particles,  or  molecules,  which  are  in 
constant  motion.  The  readiness  with  which  the  molecules 
of  gases  move  about  enables  them  to  mix  together.  The 
odour  of  musk,  which  is  due  to  the  escape  of  particles  of  the 
body  in  the  gaseous  state,  will  pervade  a  whole  room,  but  the 
finest  weighing  would  probably  not  detect  any  difference  in 
the  weight  of  the  substance  from  which  the  particles  had 
escaped  to  mingle  with  the  air  of  the  room. 

Definite  Bodies.  — The  matter  of  all  bodies  is  made  up  of 
molecules.  When  the  molecules  of  a  body  are  all  alike,  it  is 
said  to  have  a  definite  composition.  Such  bodies  have  well- 
defined  properties  by  which  they  can  be  recognised.  They 
may  be  formed  in  different  ways,  but  they  are  always  the 
same  in  composition.  It  must  not,  however,  be  assumed  that 
every  well-known  body  is  of  this  description ;  for  in  many 
cases  further  examination  brings  out  the  fact  that  different 
kinds  of  molecules  are  mixed  together  in  the  same  body. 
Such  bodies  are  complex  in  character,  and  it  is  part  of  the 
chemist's  work  to  separate  the  molecules  of  the  same  kind, 
and  recognise  the  definite  bodies  to  which  they  belong.  Com- 
mercial specimens  of  definite  bodies  are  seldom  perfectly  pure, 
but  the  adulterating  substances  are  not  usually  present  in 
sufficient  quantity  to  interfere  with  their  ordinary  properties, 
or  with  experiments  in  which  they  are  used. 

Red  Oxide  Of  Mercury.— This  well-known  substance  has 
a  definite  composition.  In  its  common  form  it  is  a  heavy, 
red,  crystalline  powder,  which  on  close  examination  is  found 


PEOPEKTIES  OF  MATTEK 


39 


to  be  quite  homogeneous.  Its  smallest  visible  particles  are 
all  alike. 

EXP.  22. — Weigh  2  grams  of  red  oxide  of  mercury  and  introduce 
it,  by  means  of  a  gutter  of  glazed  paper,  into  a  piece  of  combustion- 
tube  closed  at  one  end  and  about  5  inches  long.  Fit  the  open  end 
of  the  tube  with  a  rubber  bung  and  delivery-tube,  and  fix  the  tube 
in  a  clip  for  heating  in  the  Bunsen  flame.  Pass  the  end  of  the 
deli  very -tube  under  water  in  a  shallow  trough,  and  place  over  it 
an  inverted  gas  cylinder  of  about  100  c.c.  capacity,  previously  filled 
with  water.  Heat 
the  part  of  the 
combustion  -  tube 
which  contains  the 
red  oxide  in  the 
Bunsen  flame 
until  the  whole  of 
the  powder  dis- 
appears, or  until  a 
very  small  residue 
of  impurity  only 
is  left.  The  ar- 
rangement of  the  FIG.  14. 
apparatus  is 

shown  in  Fig.  14.  When  bubbles  of  gas  cease  to  escape  from  the 
end  of  the  delivery-tube,  remove  it  from  the  water  by  taking  the 
clip  and  tubes  away  bodily.  When  the  combustion-tube  is  cold, 
take  out  the  bung  and  remove  the  whole  of  the  mercury  which  has 
collected  on  the  sides  near  the  open  end.  If  necessary,  use  a  little 
water  to  bring  away  the  last  particles  of  metal.  A  glass  rod  tipped 
with  a  short  piece  of  rubber  tube  is  useful.  Dry  the  mercury  by 
rolling  it  about  in  a  small  tray,  made  by  turning  up  the  edges  of  a 
piece  of  blotting-paper.  Weigh  the  globule  of  metal. 

The  difference  between  the  weight  of  the  red  oxide  and 
that  of  the  metal  obtained  from  it  gives  the  weight  of  the  gas 
driven  off  during  the  heating.  On  testing  the  gas  in  the 
cylinder  with  a  glowing  splint  it  is  found  to  be  oxygen. 


EXAMPLE. — Weight  of  oxide 

,,         mercury 
„         oxygen  . 


2-000  grams 
1-846      „ 
0-147  gram 


If  a  number  of  such  experiments  were  made,  using  different 
weights  of  the  same  red  oxide,  it  would  be  proved  on  com- 
parison that  the  relation  between  the  weights  of  mercury  and 


40  METALLURGICAL  CHEMISTRY 

oxygen  obtained  from  the  oxide  is  always  the  same.  This  is 
one  way  of  deciding  whether  a  given  body  has  a  definite  com- 
position or  not.  It  may  then  be  taken  for  granted  that  red 
oxide  of  mercury  is  a  definite  body,  and  that  a  weighed 
quantity  of  it  is  made  up  of  an  enormous  number  of  exactly 
similar  molecules.  The  observed  effect  of  heat  upon  it  is  very 
clear.  It  is  split  up  into  two  other  bodies,  which  are  also  as 
definite  in  character  as  the  red  oxide  itself,  and  each  of  these 
bodies  is  made  up  of  its  own  particular  kind  of  molecules. 
Where  do  these  new  molecules  come  from  ?  They  must  be 
produced  in  some  way  from  the  disappearing  molecules  of  the 
oxide.  Careful  attention  to  the  experiment  will  make  it 
evident  that  the  action  is  a  gradual  one,  and  that  the  mercury 
and  oxygen  appear  in  proportion  as  the  oxide  disappears. 
Now,  it  should  not  be  difficult  to  picture  the  individual  mole- 
cules of  the  oxide  as  being  acted  upon  by  the  heat,  and 
divided  up  into  particles  of  mercury  and  oxygen,  which  then 
form  the  molecules  of  these  new  bodies.  Further,  these 
bodies  have  both  been  made  the  subject  of  the  most  exhaus- 
tive examination,  but  no  one  has  yet  been  able  to  divide  up 
their  molecules  so  as  to  obtain  molecules  of  other  new  bodies 
from  them.  Therefore,  in  the  light  of  present  knowledge, 
the  matter  of  either  mercury  or  oxygen  is  as  simple  as  it  can 
possibly  be. 

Clearly,  then,  definite  bodies,  of  which  red  oxide  of  mercury, 
mercury,  and  oxygen  are  examples,  belong  to  two  distinct 
types :  (1)  Those  from  which  only  one  kind  of  matter  can 
be  obtained  by  any  known  process  are  called  Elements  ; 
(2)  those  from  which  two  or  more  distinct  kinds  of  matter  can 
be  obtained  are  called  Compounds. 

The  number  of  elements  is  very  small  compared  with  the 
number  of  compounds  formed  from  them ;  but  the  beginner 
will  easily  understand  the  possibility  of  a  very  large  number 
of  compounds  being  formed  from  a  small  number  of  elements, 
if  he  compares  it  with  that  of  the  formation  of  the  words  of  a 
language  from  the  letters  of  its  alphabet. 


PROPERTIES  OF  MATTER  41 

Matter,  then,  as  we  know  it,  is  made  up  of  elements  and 
compounds ;  and  as  the  result  of  a  stupendous  amount  of  work, 
done  for  the  most  part  during  the  last  150  years,  very  definite 
statements  can  be  made.  Some  of  them  are  so  far-reaching 
that  they  are  dignified  by  the  name  of  laws.  They  will  be 
dealt  with  in  due  course.  The  following  definitions  will  be 
found  useful : 

An  element  is  a  body  from  which  only  one  kind  of 
matter  can  be  obtained.  A  compound  contains  two 
or  more  elements  in  chemical  union.  Bodies  which  do 
not  come  within  these  definitions  may  be  regarded  as  more  or 
less  intimate  mixtures  of  elements  or  compounds,  or  of  both 
elements  and  compounds,  as  the  case  may  be.  They  are 
complex  in  composition,  for  their  molecules  are  not  all  alike. 
There  are  numerous  examples  of  such  bodies. 

Now,  it  is  quite  justifiable  to  inquire  whether  the  molecules 
of  elements  are  the  smallest  possible  particles  of  these  bodies. 
Much  information  has  been  accumulated  in  answer  to  this 
question.  For  example,  are  the  particles  of  mercury  and 
oxygen  which  leave  the  red  oxide  during  its  decomposition 
really  molecules,  or  are  they  still  smaller  pieces  of  matter 
which  run  together  to  form  the  molecules  of  these  elements  in 
the  free  state  1  Some  very  concise  statements  can  be  made  at 
this  point,  and  the  beginner  may  accept  them  on  the  under- 
standing that  he  will  gradually  become  acquainted  with  the 
experimental  evidence  upon  which  they  are  based. 

Atoms. — The  molecule  of  oxygen  is  made  up  of  two 
smaller  and  exactly  similar  particles  of  oxygen  called  atoms. 
These  atoms  are  indivisible,  for  in  them  the  extreme  limit  of 
subdivision  has  been  reached. 

The  molecule  of  mercury  when  in  the  gaseous  state  is  not 
divisible,  but  is  the  smallest  possible  particle  of  the  metal.  It 
corresponds  to  the  atom  of  oxygen  in  this  respect,  and  is  also 
called  the  atom  of  mercury.  So  that  when  the  smallest  free 
particle  of  an  element  is  also  the  smallest  possible  particle  of 


42  METALLUKGICAL  CHEMISTRY 

that  element,  it  is  called  a  one-atom  or  monatomic  molecule. 
The  oxygen  molecule  is  a  more-than-one  atom,  or  polyatomic 
molecule.  These  remarks  apply  to  the  elements  generally. 

It  is  easy  to  imagine  the  oxygen  (Exp.  22)  escaping  from 
the  red  oxide  as  atoms,  and  these  atoms  running  together  in 
pairs  to  form  the  molecules  of  the  free  gas ;  and  the  mercury 
escaping  as  atoms,  which  do  not  run  together,  and  therefore 
form  molecules  by  themselves  when  in  the  form  of  vapour. 
Also,  in  the  oxidation  experiments  described  in  Chap.  II., 
the  oxygen  molecules  in  the  air  taking  part  in  the  change  may 
be  pictured  as  breaking  up  into  atoms,  and  joining  with  atoms 
of  the  metals  to  form  molecules  of  the  metallic  oxides,  which 
are  the  products  of  the  change. 

It  must  here  be  borne  in  mind  that  one  of  the  intrinsic 
properties  of  matter  is  its  weight,  and,  however  small  the 
particle  of  matter  may  be,  it  must  possess  this  property. 
Clearly,  then,  we  may  speak  of  the  weight  of  an  atom. 

An  atom  may  be  defined  as  the  smallest  particle  of  an 
element  which  can  exist,  or  which  can  take  part  in  the  forma- 
tion of  a  molecule  of  an  element  or  compound. 

A  molecule  is  the  smallest  particle  of  an  element  or  com- 
pound which  can  have  a  free  existence.  The  masses  of  all 
bodies  are  made  up  of  molecules.  The  molecules  of  the 
vapours  of  some  metals,  such  as  mercury,  zinc,  and  cadmium, 
are  monatomic.  As  soon  as  the  possibility  of  the  existence  of 
atoms  became  clear  in  the  minds  of  the  originators  of  the 
theory,  the  want  of  some  method  of  representing  them  by 
signs  or  symbols  was  felt.  At  first  these  symbols  were  some- 
what fanciful,  but  now  the  system  has  been  reduced  to  a  very 
simple  form,  in  which  the  symbol  representing  the  atom  of  an 
element  consists  of  one  or  two  letters  taken  either  from  its 
common  or  from  its  scientific  name.  Thus  the  symbol  for 
oxygen  is  0  ;  for  hydrogen,  H ;  for  copper,  Cu  (Cuprum) ;  for 
lead,  Pb  (Plumbum).  But  if  0  represents  the  atom  of  oxygen 
it  also  represents  its  weight ;  for  the  weight  of  an  atom  is 


PKOPEKTIES  OF  MATTEB  43 

intimately  connected  with  the  atom  itself.  Very  little  can  be 
said  about  the  absolute  weight  of  an  atom,  but  it  is  quite 
possible  to  compare  the  different  elements  with  regard  to  the 
relative  weights  of  their  atoms.  Hydrogen,  which  is  the 
lightest  body  known,  has  been  selected  for  the  standard  atomic 
weight,  and  the  other  elements  are  compared  with  it.  An 
account  of  the  experimental  evidence  upon  which  this  rests 
will  be  found  in  Chap.  VIIL,  and  tables  containing  the 
common  elements,  with  their  symbols  and  atomic  weights,  are 
given  on  pp.  46-7. 

The  symbolic  representation  of  atoms  only  requires  a  little 
extension  to  adapt  it  to  the  representation  of  molecules.  If 
it  is  assumed  that  both  the  number  and  the  nature  of  the 
atoms  in  a  given  molecule  are  known,  no  difficulty  presents 
itself ;  but  if  the  nature  only  of  the  atoms  in  the  molecule  is 
known,  the  best  plan  is  to  assume  that  their  number  is  as 
small  as  possible  consistent  with  the  composition  of  the  mass 
of  the  body. 

The  molecule  of  oxygen  is  said  to  contain  two  atoms  of  the 
element ;  it  could,  therefore,  Ue  r  presented  by  OO,  or  by 
20,  but  it  is  gene  ally  represented  by  O2,  and  this  combination 
of  symbols  is  called  the  formula  of  the  oxygen  molecule.  The 
representation  of  molecules  by  definite  formulae  furnishes  a 
good  example  of  the  application  of  theory  to  the  explanation 
of  the  results  of  experiment,  and  a  short  descript'on  of  the 
nature  cf  the  experimental  evidence  upon  which  themole2ular 
formula  of  oxygen  is  based  will  serve  a  useful  purpose  in  making 
it  clear  thaL  formulae  generally  depend  upon  such  evidence. 

A  general  statement  which  is  strongly  supported  by  experi- 
mental evidence  can  be  made  about  gaseous  bodies.  It  is 
known  as  the  Law  of  Avogadro,  and  is  expressed  thus  : 

Equal  volumes  of  all  gases  under  the  same  condi- 
tions of  temperature  and  pressure  contain  the  same 
number  of  molecules. 

Thus  equal  volumes  of  oxygen  and  hydrogen  contain  the 


44  METALLURGICAL  CHEMISTRY 

same  number  of  molecules,  and  double  the  volume  of  either 
gas  would  contain  double  the  number  of  molecules.  When 
two  volumes  of  hydrogen  are  mixed  with  one  volume  of 
oxygen  at  a  temperature  above  the  boiling-point  of  water, 
simple  mixture  takes  place,  and  three  volumes  of  gas  are 
obtained.  If  the  mixture  is  ignited,  the  gases  combine  and 
s  earn  is  formed,  whijh  on  cooling  to  the  original  tempe-ature 
of  the  mixture  occupies  exactly  two  volumes.  Thus  the 
volume  of  the  steam  is  double  that  of  the  oxygen  which  helps 
to  form  it,  so  that  the  matter  of  which  the  oxygen  molecules 
are  composed  is  distributed  into  twice  as  many  molecules  of 
steam.  That  is,  each  oxygen  molecule  is  divided  into  two 
parts  during  the  chemical  change  by  which  the  hydrogen  and 
oxygen  molecules  are  converted  into  molecules  of  water.  If, 
then,  the  law  of  Avogadro  is  true,  oxygen  molecules  must 
contain  at  least  two  atoms  of  the  element,  and  as  no  change 
is  known  in  which  they  appear  to  be  divided  into  more  than 
two  pirts,  the  inference  that  oxygen  molecules  contain  two 
atoms  is  justified. 

The  weight  of  a  given  volume  of  any  gas  compared  with 
the  weight  of  an  equal  volume  of  hydrogen,  which  is  taken  as 
the  unit  weight,  is  called  its  density.  If,  then,  it  is  assumed 
that  a  molecule  of  hydrogen  is  made  up  of  two  atoms,  each 
of  unit  weight,  the  molecular  weight  of  the  gas  is  double  its 
atomic  weight.  By  the  application  of  Avogadro's  law  it 
follows  that  the  molecular  weights  of  gases  are  proportional 
to  their  densities.  So  that  if  the  molecular  weight  of  hydro- 
gen is  double  its  density,  the  molecular  weights  of  other  gases 
are  obtained  by  doubling  their  densities  compared  with  that 
of  hydrogen.  Thus  if  an  element  is  a  gas,  or  can  be  con- 
verted into  gas  at  a  workable  temperature,  and  its  atomic 
weight  is  known,  the  number  of  atoms  in  its  molecule  is 
determined  by  comparing  the  weight  of  a  given  volume  of  the 
gas  with  that  of  an  equal  volume  of  hydrogen  under  the  same 
conditions  of  temperature  and  pressure.  This  method  of 
determining  molecular  weights  can  be  applied  to  all  elements 


PROPERTIES  OF  MATTER  45 

and  compounds  obtainable  in  the  gaseous  state.  The  assump- 
tion that  the  molecule  of  hydrogen  contains  two  atoms  is 
strongly  supported  by  evidence  similar  to  that  given  above  in 
the  case  of  the  oxygen  molecule. 

Symbols  and  formulae  are  used  primarily  to  represent 
atoms  and  molecules,  but  there  is  another  way  of  using  them 
which  is  of  great  practical  importance.  It  depends  upon 
reasoning  such  as  that  which  follows  : 

A  given  volume  of  oxygen  is  made  up  of  an  enormous  but 
definite  number  of  molecules,  and  if  n  represent  this  number, 
the  weight  of  the  given  volume  is  represented  by  7i02.  The 
atomic  weight  of  oxygen  is  0=16,  and  its  molecular  weight 
is  O2  =  32.  Now,  it  is  easy  to  see  that  n  may  be  such  a  number 
that  nQ%  represents  a  definite  weight  of  the  gas,  say  32  grams  ; 
so  that  n0.2  might  be  used  to  represent  a  molecular  weight  of 
oxygen  in  grams,  or  any  other  standard  weights.  This  is 
done,  but  it  is  usual  to  omit  the  n.  Thus  the  formula  02  is 
used  in  two  senses:  (1)  to  describe  a  single  molecule  of 
oxygen  ;  (2)  to  denote  a  definite  weight  of  the  free  gas,  as 
32  grams  32  pounds,  etc.  The  number  32  always  appears  in 
any  statement  which  concerns  the  weight  of  a  molecule  of 
oxygen,  or  any  number  of  the  same.  The  symbols  and 
formulae  of  bodies  in  general  are  used  in  the  same  way.  .'  /.-/ 

Elements. — These  bodies  have  characteristic  physical  and 
chemical  properties,  by  which  they  can  be  recognised  and 
distinguished  from  each  other,  either  by  general  inspection  or 
by  chemical  tests.  They  are  divided  into  two  main  groups. 

Metals. — The  larger  group  contains  the  metals,  several  of 
which  are  so  common,  and  have  such  well-defined  properties, 
that  everyone  is  more  or  less  acquainted  with  them.  The 
metals  have  two  characteristic  properties  which  serve  to 
distinguish  them  from  other  elements  :  (1)  The  smooth  clean 
surface  of  a  metal  has  the  peculiar  appearance  known  as 
metallic  lustre  ;  (2)  the  metals  are  good  conductors  of  heat 
and  electricity.  Most  of  the  common  metals  are  very  heavy 


46 


METALLURGICAL  CHEMISTRY 


bodies,  but  they  vary  very  much  among  themselves  with 
regard  to  this  property.  Platinum,  the  heaviest  of  the  com- 
mon metals,  is  twenty-five  times  heavier  than  potassium,  one 
of  the  lightest  of  them. 

Non-metals. — The  smaller  group  includes  a  number  of  ele- 
ments whose  properties  are  so  opposite  to  those  of  the  metals 
that  they  are  known  as  non-metals.  They  have  no  common 
physical  characteristics,  but  their  chemical  properties  are  so 
marked,  that  they  are  taken  advantage  of  in  classifying  them. 
This  group  includes  such  physically  different  bodies  as  sulphur, 
carbon,  and  oxygen.  A  few  of  the  elements  occupy  the  border- 
line between  the  well-defined  metals  and  non-metals,  and  seem 
to  partake  of  the  characters  of  both.  The  term  metalloids  is 
applied  to  these,  but  they  need  not  be  considered  at  present. 

Some  elements  occur  in  such  enormous  quantities  as  to  be 
excessively  common ;  others  in  much  smaller  quantities,  but 
so  widely  distributed  as  to  be  still  common ;  others  again  in 
small  quantities,  and  of  limited  distribution.  These  are  called 
rare  elements,  and,  although  interesting  from  the  chemist's 
point  of  view,  are  of  little  known  practical  value.  A  list  of 
them  is  given  in  the  Appendix. 

The  following  tables  contain  the  names,  symbols,  and  atomic 
weights  of  the  elements  which  are  usually  found,  either  in  the 
free  state  or  as  compounds,  on  the  laboratory  shelves.  Full 
lists  will  be  found  in  the  Appendix. 

COMMON  METALS. 


2 

ii 

i 
? 

11 

1 

J"*  JS 

Name. 

fO 

3 

Name. 

^- 

£ 

|S 

Name. 

#3 

a 

i 

g 

>. 

>» 

e 

p 

OD 

^ 

OQ 

•4* 

Aluminium 
Antimony 

Al 
Sb 

27-0 
120-0 

Copper             Cu 
Gold                 Au 

63-0 
197-0 

Nickel           Ni 
Potassium     K 

58-5 
39-0 

Barium          Ba  !  137'0 

Iron 

Fe 

56-0 

Platinum       Pt 

194-0 

Bismuth        Bi     208  -0 

Lead               i  Pb 

207-0 

Silver           1  Ag 

108-0 

Cadmium 

Cd 

112-0 

Magnesium      Mg 

24-0 

Sodium          Na 

23-0 

Calcium 

Ca 

40-0 

Manganese     ;  Mn 

55-0 

Strontium 

Sr 

87-0 

Chromium 
Cobalt 

Cr 
Co 

52-5 

58-5 

Mercury 
Molybdenum 

Hg 
Mo 

200-0 
96-0 

Tin 
Zinc 

Sn 
Zn 

118-0 
65-0 

PEOPEETIES  OF  MATTEE 


47 


COMMON  NON-METALS. 


Name. 

i 

14 

g_bc    ; 

Name. 

2 

II 

Name. 

| 

.2  *" 
11 

^> 

-*•*  ^ 

t» 

ijj  jP 

>» 

x 

OQ 

DO 

^^ 

Arsenic 

As 

75-0 

Fluorine 

F 

19-0 

Oxygen 

0 

16-0 

Boron 

B 

11-0 

Hydrogen 

H 

1-0 

Phosphorus 

p 

31-0 

Bromine 

Br 

80-0 

Iodine 

i 

127-0 

Sulphur 

s 

32-0 

Carbon 

C 

12-0    Nitrogen 

N 

14-0 

Silicon 

Si 

28-0 

Chlorine 

Cl 

35-5 

Mercury  is  the  only  common  metal  which  is  liquid  at  atmo- 
spheric temperatures.  Arsenic  is  sometimes  classed  with  the 
metals,  but  its  chemical  properties  are  more  like  those  of  the 
non-metals.  The  atomic  weights  given  are  only  approxi- 
mate, but  they  will  be  found  sufficiently  accurate  for  ordinary 
work. 

Compounds. — These  bodies  are  just  as  definite  in  their 
chemical  characters  and  in  their  constancy  of  composition  as 
the  elements  themselves.  This  cannot  be  too  strongly  urged. 
In  fact,  the  generalizations  arrived  at  as  the  result  of  much 
labour  by  many  men  are  so  all-embracing  that  they  are  put 
forward  with  much  confidence,  and  are  looked  upon  as  thor- 
oughly reliable  guides.  They  are  known  as  the  laws  of 
chemical  combination.  The  first  one  may  be  stated  here,  the 
others  will  be  given  later. 

The  first  law  of  chemical  combination  is  as  follows  : 

A  given  compound  always  contains  the  same  ele- 
ments, and  in  the  same  proportions. 

Verifications  of  this  law  will  be  constantly  presenting 
themselves,  and  the  student  should  find  no  difficulty  in 
finally  accepting  it.  Two  examples  may  now  be  given.  Red 
oxide  of  mercury  always  contains  mercury  and  oxygen  in  the 
proportion  of  12'5  parts  of  mercury  to  1  part  of  oxygen  by 
weight.  Water  always  contains  oxygen  and  hydrogen  in 
the  proportion  of  1  part  of  oxygen  to  2  parts  of  hydrogen  by 


48  METALLURGICAL  CHEMISTRY 

volume  (see  Exp.  19).  Or,  since  oxygen  is  sixteen  times 
heavier  than  hydrogen,  water  contains  8  parts  of  the  former 
to  1  part  of  the  latter  by  weight.  This  will  be  confirmed 
later. 

Formulce  of  Compounds.  —  Now,  in  Exp.  19  it  is  proved  that 
2  grams  of  red  oxide  of  mercury  contains  1-846  grams  of 
mercury  and  0*147  gram  of  oxygen.  Assuming  that  the  weight 
of  each  element  is  made  up  of  a  definite  number  of  atoms  of 
the  element,  the  ratio  between  these  numbers  can  be  found  by 
dividing  each  weight  by  the  atomic  weight  of  the  element  to 
which  it  belongs.  Thus  : 

=  0.0092  =  0-0092. 


The  ratio  is  1  :  1  —  that  is,  the  two  weights  contain  the 
same  number  of  atoms.  So  that  they  must  leave  the  oxide 
molecules  in  pairs,  and  are,  therefore,  present  in  the  mole- 
cules in  pairs.  Clearly,  then,  the  simplest  way  to  represent 
a  molecule  of  the  oxide  is  to  put  the  symbols  together  thus— 
HgO. 

In  the  decomposition  of  water  by  electrolysis  it  is  proved 
that  the  constituent  gases  are  liberated  in  the  proportion  of 
2  volumes  of  hydrogen  to  1  volume  of  oxygen.  Also, 
by  Avogadro's  law,  the  double  volume  of  hydrogen  contains 
twice  as  many  molecules  as  the  single  volume  of  oxygen  ;  and, 
since  all  these  molecules  are  diatomic,  it  is  probable  that  the 
hydrogen  and  oxygen  atoms  leave  the  water  molecules  in  the 
proportion  of  2  :  1.  Therefore  the  simplest  way  of  expressing 
the  water  molecule  is  to  write  the  symbols  together  thus  — 
H20.  And  there  is  much  evidence  in  favour  of  this  as  the 
actual  composition  of  a  molecule  of  water.  These  complex 
symbols  are  called  formulce.  Thus,  HgO  and  H2O  are  the 
formulae  for  red  oxide  of  mercury  and  water.  Next  consider 
Exp.  2.  It  is  not  a  case  of  the  indifferent  absorption  of 
oxygen  into  lead  as  a  mass,  but  the  individual  atoms  of 
oxygen  become  associated  with  the  individual  atoms  of  lead, 


PEOPEETIES  OF  MATTER  49 

and  form  molecules  of  oxide  of  lead.  At  any  intermediate 
stage  the  mass  consists  of  two  parts  :  (a)  oxide  of  lead  ; 
(b)  unchanged  metal.  It  is  either  one  or  the  other  ;  there  is 
no  intermediate  change.  In  this  transformation  of  lead  it  is 
found  that  207  parts  by  weight  of  lead  and  16  parts  by  weight 
of  oxygen  combine  to  form  223  parts  of  the  oxide.  But  as 
these  are  the  atomic  weights  of  the  two  elements,  it  appears 
that  the  atoms  of  lead  and  oxygen  run  together  in  pairs  to 
form  the  molecules  of  the  oxide.  Therefore,  such  a  statement 
as  the  following  may  be  made  : 

n  atoms  of  lead  (nPlo)  and  n  atoms  of  oxygen  (nO)  furnish 
n  molecules  of  oxide  (?iPbO).  Or,  replacing  and  by  +  ,  and 
furnish  by  =,  and  taking  out  the  n,  which  is  understood,  it 
may  be  written  : 

=  PbO. 


This  statement  is,  as  it  stands,  a  chemical  equation,  and 
expresses  the  chemical  change  under  discussion.  But,  as  the 
oxygen  of  the  air  is  in  the  form  of  molecules,  it  is  usual  to 
write  the  equation  thus  : 


Here  2PbO  means  2[PbO],  that  is  two  molecules  of  the  oxide, 
and  the  same  mode  of  expressing  a  number  of  molecules  is 
always  used. 

The  decomposition  of  red  oxide  of  mercury  is  expressed  as 
follows  : 


And  the  equation  reads  thus  :  2  molecules  of  red  oxide  of 
mercury  furnish  2  molecules  of  mercury  and  1  molecule  of 
oxygen.  The  beginner  will  probably  have  a  little  difficulty 
in  fully  understanding  the  meaning  of  symbols,  formula?,  and 
equations,  but  as  he  progresses  with  the  practical  work,  and 
becomes  familiar  with  them  by  use,  the  difficulty  will  vanish. 

Impurities.  —  Ordinary  specimens  of   elements  and  com- 

4 


50  METALLURGICAL  CHEMISTRY 

pounds  usually  contain  small  quantities  of  other  bodies  which 
are  regarded  as  impurities,  but  these  do  not  as  a  rule  seriously 
affect  the  characteristic  properties  of  the  bodies  under  con- 
sideration. Almost  pure  specimens  can,  however,  be  obtained 
if  necessary.  For  example,  gold  containing  less  than  1  part 
of  impurity  in  10,000  parts  of  the  metal  is  to  be  obtained,  but 
it  costs  about  three  times  as  much  as  ordinary  fine  gold, 
which  contains  about  3  parts  of  impurity  in  1,000  parts  of  the 
metal.  Zinc  and  copper  containing  less  than  1  part  of  im- 
purity in  2,000  parts  of  the  metals  are  easily  obtained,  and 
many  compounds  are  produced  in  a  state  of  great  purity. 

Complex  Bodies. — These  bodies  are  very  numerous  and 
varied.  Many  of  them  are  common,  and  have  some  well- 
defined  properties ;  but  they  lack  that  absolute  constancy  of 
composition  which  is  characteristic  of  elements  and  compounds, 
The  atmosphere  may  be  taken  as  an  example.  Pure  mountain 
air  contains  at  least  four  substances  of  definite  composition  : 
two  elements,  oxygen  and  nitrogen,  and  two  compounds, 
water  vapour  and  carbon  dioxide.  The  proportion  of  oxygen 
to  nitrogen  is  very  nearly  constant,  but  the  evidence  is  all  in 
favour  of  a  simple  mixture  of  these  gases  in  air.  The  pro- 
portion of  carbon  dioxide  varies  slightly,  and  that  of  water 
vapour  considerably,  from  time  to  time.  Granite  is  found  to 
contain  at  least  three  compounds,  which  on  examination  of 
the  polished  surface  of  the  stone  are  seen  to  be  simply  mixed 
together. 

Mixtures  are  often  made  for  practical  purposes,  and  gun- 
powder, which  contains  saltpetre,  sulphur,  and  charcoal,  is  a 
good  example.  Such  bodies  are  called  mechanical  mix- 
tures, and  when  examined  under  the  microscope  the  various 
constituents  are  readily  recognised. 

The  constituents  of  mixtures  can  usually  be  separated  from 
each  other  without  resorting  to  chemical  changes.  Such 
separation  can  be  effected  by  mechanical  means.  The 
removal  of  particles  of  iron  from  brass-dust,  and  of  oxide  of 


PROPERTIES  OF  MATTER  51 

iron  from  magnetic  sands  by  magnets,  are  examples.  If 
gunpowder  is  shaken  up  with  hot  water,  the  saltpetre  in  it 
dissolves,  and  may  be  completely  separated  by  filtration  and 
washing;  the  contained  sulphur  is  just  as  easily  dissolved 
from  the  dried  residue  by  digesting  it  with  carbon  bisulphide; 
and  the  charcoal  is  left  as  a  black  residue.  On  evaporating 
the  solutions  of  the  saltpetre  and  sulphur  respectively,  the 
solids  are  obtained  in  the  same  condition  as  they  were  when 
used  to  make  the  original  mixture.  Complete  separation  is 
thus  effected  without  any  permanent  change  in  the  con 
stituents  of  the  mixture. 

Theory  of  Atoms  and  Molecules.— In  the  above  state- 
ments of  facts  about  matter  in  general,  and  in  the  attempt  to 
explain  them,  the  broad  aspects  of  the  atomic  theory  have 
been  used.  The  original  idea  of  atoms  is  by  no  means  a  new 
one,  but  it  was  introduced  in  its  modern  dress  by  Dalton,  of 
Manchester,  in  1808,  who  used  it  most  successfully  in  explain- 
ing some  facts  of  chemical  combination  which  he  had  dis- 
covered. Since  his  time  other  considerations  have  necessitated 
an  extension  of  the  original  ideas,  so  as  to  distinguish  between 
atoms  and  molecules.  It  is  in  this  extended  sense  that 
Dalton's  atomic  theory  is  used  by  the  modern  chemist. 

When  Dalton  adopted  the  idea  that  matter  is  made  up  of 
small  particles  (atoms),  he  assumed  that  they  are  all  alike  in 
a  given  body  of  constant  composition.  If  this  is  not  so,  then 
it  should  be  possible  to  produce  two  different  specimens  of 
the  same  compound  from  its  elements  by  a  given  chemical 
change.  For  if  the  smallest  particles  of  the  elements  differ  in 
size  in  the  quantities  used,  the  particles  of  the  compound 
should  differ  also.  But  no  such  variations  in  the  properties 
of  different  specimens  of  the  same  compound  are  known ;  it  is, 
therefore,  a  fair  inference  that  the  atoms  of  an  element  are  the 
exact  counterparts  of  each  other.  It  also  follows  that  if  the 
atoms  of  elements  are  alike,  the  molecules  of  compounds 
formed  from  them  must  be  alike.  These  assumptions  enable 

4—2 


52  METALLUKGICAL  CHEMISTEY 

us  to  account  in  a  simple  way  for  the  constancy  in  the  com- 
position of  elements  and  compounds. 

It  has  already  been  noticed  that  there  are  two  distinct 
compounds  of  copper  aud  oxygen,  and  it  is  well  known  that 
the  weights  of  copper  and  oxygen  in  these  compounds  are  in 
the  proportions  of  63  :  8  and  63  :  8x2,  and  in  no  others. 
It  was  evidence  of  this  kind,  obtained  from  experiments  with 
some  of  the  compounds  of  carbon,  that  led  Dalton  to  the 
discovery  of  the  law  of  chemical  combination  in  multiple  propor- 
tions, which  may  be  stated  thus  : 

When  one  element  unites  with  another  element  in 
more  than  one  proportion,  the  higher  proportions 
are  simple  multiples  of  the  first  proportion. 

This  law  holds  rigorously  when  two  elements  combine  in 
different  proportions  to  form  two  or  more  distinct  compounds; 
and  it  seems  very  probable  that  the  molecules  of  the  com- 
pounds containing  the  higher  proportions  are  formed  from  the 
molecules  of  those  containing  the  lower  proportion,  by  the 
addition  of  like  particles  of  the  other  element  to  them.  This 
was  the  view  that  Dalton  took,  and  it  enabled  him  to  use  the 
atomic  hypothesis  in  explaining  the  facts  he  had  discovered. 

It  is  necessary  to  distinguish  clearly  between  a  series  of 
facts  and  the  efforts  of  the  mind  to  explain  them.  A  theory 
is  a  mind  picture,  which  keeps  our  work  out  of  the  mechanical 
rut,  and  raises  it  from  the  region  of  mere  groping  to  that  of 
intelligent  effort.  A  good  theory  should  not  only  be  able  to 
explain  what  has  happened,  but  also  predict  what  may  happen 
under  certain  conditions.  And  the  theory  of  atoms  and 
molecules  has  done  so  much  in  this  direction,  that  it  is 
accepted  with  the  utmost  confidence,  and  is  now  regarded 
almost  as  a  proved  truth.  The  student  will,  therefore,  find 
good  mental  training  in  looking  for  evidence  in  its  support, 
and  in  using  it  to  explain  observed  facts. 


PROPERTIES  OF  MATTER  53 

SUMMARY. 

Matter  is  the  substance  of  a  body.  The  term  "  body  "  is  used 
to  denote  a  limited  quantity  of  matter  which  has  a  definite 
volume,  however  great  or  small  it  may  be.  On  the  other  hand, 
substance  is  the  general  term  for  the  matter  of  a  body,  with- 
out reference  to  any  particular  quantity  of  it.  Bodies  can  be 
divided  into  a  number  of  small  parts,  but  there  is  a  limit  to 
their  division.  Molecules  are  the  smallest  possible  free  par- 
ticles of  matter  of  any  kind.  Molecules  under  certain  con- 
ditions yield  still  smaller  particles,  called  atoms,  but  these 
cannot  have  a  free  existence  in  the  popular  sense.  The 
definitions  of  atom  and  molecule  should  be  carefully  noted, 
and  looseness  in  their  use  avoided.  They  are  not  interchange- 
able. An  atom  is  indivisible,  and  if  the  smallest  free  particles 
of  an  element  are  indivisible,  they  are  monatomic  molecules 
as  well  as  atoms.  The  term  "atom"  is  to  be  used  only  in  con- 
nection with  elements,  which  are  bodies  containing  but  one 
kind  of  matter.  Atoms  are  the  constituent  parts  of  mole- 
cules, and  the  latter  are  the  smallest  particles  of  compounds 
and  of  free  elements. 

There  is  reason  to  believe  that  some  compounds  are,  when 
in  dilute  solution,  split  up  or  dissociated  into  their  elements, 
the  atoms  of  which  are  more  or  less  free.  These  are  the 
wandering  atoms  or  ion-i  of  the  electro-metallurgist.  Also  it 
is  possible  to  obtain  some  elements  in  a  simpler  state  of 
division  than  that  in  which  they  exist  in  the  ordinary  way,  by 
exposing  them  to  a  very  high  temperature.  But  in  both 
cases  the  atoms  come  together  again  when  the  separating 
cause  is  removed.  Atoms  are  represented  by  symbols,  mole- 
cules by  formula).  To  indicate  the  number  of  similar  molecules, 
a  figure  is  written  in  front  of  the  formula  of  the  molecule,  as 
2H2,  3PbO. 

A  chemical  equation  is  the  same  as  any  other  equation, 
with  respect  to  the  equality  of  its  two  sides.  But  this  equality 
refers  only  to  the  number  and  quality  of  the  atoms  on  the 


54  METALLURGICAL  CHEMISTRY 

two  sides,  which  must  be  the  same,  however  varied  their 
redistribution  in  the  molecules  may  have  been.  Thus  in  the 
equation 

=  2PbO, 


there  are  two  atoms  of  lead  and  two  atoms  of  oxygen  on  both 
sides,  but  the  atoms  of  the  molecules  of  the  elements  on  the 
one  side  have  been  redistributed  into  molecules  of  oxide  of 
lead  on  the  other. 

A  theory  should  always  be  accepted  with  an  open  mind, 
and  judged  according  to  its  ability  to  explain  observed  facts. 
If  it  does  this  satisfactorily,  as  is  the  case  with  the  atomic 
theory,  it  may  be  used  freely. 

NOTE.  —  The  investigation  of  radio-active  bodies,  which  received 
such  an  impetus  by  the  discovery  of  radium,  seems  to  point  to  the 
slow  breaking  down  of  atoms  into  still  smaller  particles.  This 
can,  however,  be  only  casually  referred  to  in  an  elementary  work 
until  the  nature  of  these  particles  is  better  understood,  and  its 
influence  upon  our  present  idea  of  atoms-  more  settled. 

QUESTIONS. 

1.  Give  a    definition   of   matter.      Give   examples  of   the 
division  of  bodies  into  small  particles. 

2.  What  is  a   refinite  body  ?     Illustrate  your  answer  by 
reference  to  water. 

3.  Describe   an   experiment   in   which  decomposition   into 
simpler  forms  of  matter  takes  place. 

4.  Define  the  terms  atom,  molecule,  element,  compound. 

5.  How  may  symbols  be  used  to  represent  the  proportions 
of  the  elements  in  a  compound  1 

6.  What  do   you  understand  by  the  terms  u  hypothesis  " 
and  "  theory  "  ? 

7.  Find   the  formulae  of  the   oxides   of   iron  and  copper, 
having  given  that  4  grams  of  iron  combine  with  1*714  of 
oxygen,  and  6  grams  of  copper  with  1-52  of  oxygen. 


CHAPTER  V 
METALS  AND  SULPHUR 

Sulphur. — This  well-known  element  is  found  associated  with 
earthy  matter  in  volcanic  districts.  It  is  present  in  tbe 
elementary  state,  and  is  called  native  sulphur.  The  crude 
material  is  carefully  heated,  which  causes  the  sulphur  bo  melt 
and  drain  away,  leaving  the  greater  part  of  the  earthy  matter 
behind.  The  substance  thus  obtained  is  still  impure,  and  is 
known  as  crude  sulphur.  To  complete  the  purification  the 
crude  material  is  put  into  a  retort  and  strongly  heated,  when 
the  sulphur  is  converted  into  vapour.  This  passes  through 
an  opening  in  the.  retort  into  a  cooling  chamber,  where  it 
condenses  to  the  liquid  state.  The  hot  liquid  collects  in  the 
bottom  of  the  chamber,  from  which  it  is  allowed  to  run  into 
wooden  moulds,  and  cast  into  cylindrical  sticks.  This  variety 
is  called  roll  SUlphUF.  If  the  chamber  into  which  the  vapour 
passes  is  sufficiently  cool  the  condensed  material  takes  the 
form  of  a  fine  powder,  which  is  called  flowers  of  sulphur. 
These  are  the  common  forms  of  commercial  sulphur,  and  also 
of  the  purest. 

A  considerable  quantity  of  sulphur  is  also  obtained  as  a 
by-product  in  certain  chemical  and  metallurgical  operations, 
so  that  we  are  not  entirely  dependent  upon  the  native  form  of 
the  element. 

Action  of  Heat  on  Sulphur.— Roll  sulphur  is  a  pale 
yellow,  brittle,  crystalline  solid,  which,  when  heated,  melts  to 
a  clear  amber-coloured  liquid  at  114'5°C.  On  continuing  the 
heating,  the  liquid  gets  much  darker  in  colour,  and  between 


56  METALLURGICAL  CHEMISTRY 

220°  and  250°  C.  is  thick  or  viscous  like  treacle  ;  above  the 
latter  temperature  it  becomes  thin  again,  but  is  still  dark- 
coloured;  at  448°  C.  it  boils,  giving  off  a  dark-red  vapour. 
On  cooling,  the  vapour  condenses  to  the  liquid  again,  which 
then  passes  through  similar  changes  to  the  above,  but  in  the 
inverse  order,  and  so  returns  to  the  solid  state.  If,  when  in 
the  thick  condition,  the  liquid  is  poured  into  cold  water,  a 
soft  solid  mass,  called  plastic  sulphur,  is  obtained. 

EXP.  23. — Break  up  some  roll  sulphur  into  small  pieces,  and  half 
fill  a  test-tube  with  it.  Heat  the  tube  in  the  Bunsen  flame,  and 
observe  the  changes  which  take  place  up  to  the  boiling-point.  Allow 
the  tube  to  cool,  and  when  the  liquid  gets  thick  pour  the  greater 
part  of  it  into  cold  water.  Note  the  change  in  colour  of  the  remain- 
ing liquid,  and  when  it  is  amber-coloured  pour  some  into  water. 
The  residue  solidifies  on  the  side  of  the  tube.  Examine  the  water- 
cooled  portions  :  the  one  is  a  stringy,  plastic  mass ;  the  other 
consists  of  hard,  brittle  lumps. 

It  may  be  inferred  that  to  get  sulphur  into  the  plastic  form 
its  temperature  must  be  raised  above  220°  C.  before  it  is 
suddenly  cooled.  Also  that  some  change  in  the  molecules 
themselves,  or  in  their  relation  to  each  other,  must  take  place 
about  that  temperature,  and  that  the  sudden  cooling  prevents 
the  reversal  of  this  change.  The  -  plastic  mass  passes  slowly 
back  to  the  hard  yellow  form  at  ordinary  temperatures,  arid 
more  rapidly  at  100°  C. 

EXP.  24. — Melt  about  ^  pound  of  sulphur  in  a  large  clay  crucible, 
and  allow  it  to  cool  slowly  until  a  crust  forms  on  the  surface. 
Break  the  crust,  and  pour  the  remaining  liquid  into  a  dish  of  water. 
The  sides  and  bottom  of  the  crucible  are  found  to  be  covered  with  a 
mass  of  long,  transparent,  needle-shaped  crystals. 

EXP.  25. — Put  about  a  gram  of  powdered  sulphur  into  a  test-tube, 
and  add  to  it  about  5  c.c.  of  carbon  bisulphide.  Shake  the  tube 
until  most  of  the  sulphur  is  dissolved ;  then  pour  the  clear  liquid 
into  a  small  evaporating  basin,  nearly  cover  the  basin  with  a  clock 
glass,  and  put  it  in  a  fume  chamber  to  evaporate.  When  some  well- 
defined  crystals  have  formed,  pour  off  the  remaining  liquid,  and 
remove  them  for  inspection. 

On  looking  at  the  crystals  with  a  strong  magnifying-glass 
or  under  a  microscope  of  low  power,  some  well-formed  speci- 


METALS  AND  SULPHITE  57 

mens  will  be  observed.  They  are  generally  of  two  shapes, 
similar  to  a,  b  (Fig.  15),  but  the  first  shape  usually  predominates. 
They  both  belong  to  the  rhombic 
system. 

The  experiments  described  above 
reveal  three  distinct  forms  of 
sulphur. 

1.  Octahedral  crystals  belonging 
to  the  rhombic  system,    which   is 
the   stable   form   found  in   native 
sulphur,  roll  sulphur,  etc. 

2.  Needle-  or  prism-shaped  crystals  belonging  to  the  mono- 
clinic  system,  which  gradually  pass  back  to  the  stable  form, 
with  evolution  of  heat.     Melting-point  120°  C. 

3.  The  plastic  form,  which  is  insoluble  in  carbon  bisulphide, 
and  gradually  passes  back  to  the  stable  form,  with  evolution 
of  heat. 

These  are  allo tropic  modifications  of  sulphur.  This 
subject  will  be  referred  to  again  in  a  later  chapter,  bub  for 
fuller  details  of  the  varieties  of  sulphur  a  larger  work  must  be 
consulted. 

Action  of  Heat  and  Air  on  Sulphur. — When  any  form 
of  sulphur  is  heated  in  air  it  melts,  and  then  takes  fire,  burn- 
ing with  a  pale  blue  flame.  Oxygen  from  the  air  unites  with 
the  sulphur  to  form  a  colourless  transparent  gas,  having  a 
powerful  odour,  which  is  usually  described  as  the  smell  of 
burning  sulphur.  It  dissolves  readily  in  water,  and  the  solu- 
tion smells  of  the  gas,  which  is  a  product  of  the  oxidation  of 
sulphur,  and  is,  therefore,  an  oxide.  It  is  proved  by  experi- 
ment to  contain  equal  weights  of  sulphur  and  oxygen.  The 
atomic  weights  of  its  elements  are  S  =  32  and  0  =  16,  so 
that  its  simplest  formula  is  S02.  This  is  also  the  true 
molecular  formula  of  the  gas,  as  indicated  by  its  vapour 
density. 

Metals   and  Sulphur. — When    the   common    metals   are 


58  METALLUKGICAL  CHEMISTRY 

heated  with  sulphur  a  very  marked  change  takes  place,  which 
is  as  well  defined  as  that  brought  about  between  the  same 
metals  and  the  oxygen  of  the  air  under  the  influence  of  heat. 
The  compounds  formed  are  metallic  sulphides,  and  include  a 
number  of  very  important  metalliferous  bodies.  Accurate 
quantitative  results  are  easily  obtained,  so  that  experi- 
ments on  the  formation  of  sulphides  give  much  useful  in- 
formation. 

Copper. — EXP.  26. — Weigh  accurately  2  grams  of  finely- divided 
copper,  either  in  the  form  of  filings  or  of  the  reduced  metal. 
Transfer  it  to  a  clean  dry  test-tube  by  means  of  a  paper  gutter. 
Weigh  roughly  1  gram  of  powdered  sulphur,  and  add  it  to  the 
copper,  well  mixing  the  two  by  carefully  shaking  the  tube.  Heat 
the  end  of  the  tube  in  the  Bunsen  flame  until  the  action  commences, 
and  when  it  has  ceased  return  the  tube  to  the  flame,  and  heat  it 
strongly  until  the  excess  of  sulphur  is  driven  well  up  the  tube. 
Allow  the  tube  to  cool  in  a  slanting  position,  so  as  to  prevent  any 
liquid  sulphur  from  running  back  on  to  the  solid  residue.  Invert 
the  cold  tube  over  a  piece  of  glazed  paper,  and  let  the  bluish  black 
residue  fall  out.  Weigh  this  residue.  Grind  it  up  in  a  mortar,  and 
look  at  the  powder  under  a  good  magnifying-glass.  There  is  no 
trace  of  either  free  copper  or  sulphur  to  be  seen ;  the  body  appears 
to  be  perfectly  homogeneous.  Mix  a  little  copper  and  sulphur 
together ;  look  at  the  mixture  under  the  glass  ;  put  it  into  a  test- 
tube  with  a  little  water,  and  well  shake  the  tube.  The  copper 
settles  to  the  bottom,  and  the  sulphur  is  suspended  in  the  water. 
No  such  separation  can  be  brought  about  with  the  residue  after 
the  mixture  has  been  heated. 

EXAMPLE. — Weight  of  residue  2-507  grams 

,,         copper  2-000      „ 

,,         sulphur         ...         ...     0*507  gram 

That  is,  2  grams  of  copper  combine  with  0'507  gram  of  sulphur  to 
form  2-507  grams  of  copper  sulphide. 

The  weight  of  the  sulphur  absorbed  by  the  copper  is 
obtained  by  difference,  but  this  is  perfectly  legitimate,  for 
practically  no  copper  escapes  under  the  conditions  of  the 
experiment ;  it  is,  therefore,  quite  clear  that  the  increase  in 
weight  is  due  to  the  absorption  of  sulphur  by  the  metal. 


METALS  AND  SULPHUR  59 

This  definite  body  with  its  constant  composition  is  copper 
sulphide. 

Now,  bearing  in  mind  the  idea  that  the  actual  weight 
of  each  element  in  the  sulphide  is  made  up  of  a  large  number 
of  atoms,  and  that  the  relative  weights  of  these  atoms  are 
Cu  =  63,  and  S  =  32,  it  is  easy  to  understand  that,  by 
dividing  the  actual  weights  by  the  relative  weights,  the 
ratio  between  the  numbers  of  the  atoms  of  copper  and  sulphur 
combining  to  form  the  molecules  of  sulphide  is  obtained. 
Thus: 

^  =  0-032  and  ^^  =  0-016. 
oo  oZi 

The  ratio  is  evidently  2:1,  and  the  simplest  formula  for  the 
compound  is  Cu2S. 

Iron. — EXP.  27. — Repeat  the  last  experiment,  using  2  grams  of 
fine  iron  filings  and  1*5  grams  of  sulphur.  After  heating,  the 
residue  sticks  to  the  sides  somewhat,  and  it  is  necessary  to  break 
the  tube  and  remove  the  pieces  of  glass  from  the  residue  before 
weighing  it. 

EXAMPLE. — Weight  of  residue          3-145  grams 

,,          iron 2-000      ,, 

,,          sulphur         1-145      ,, 

That  is,  2  grams  of  iron  combine  with  1*145  grams  of  sulphur  to 
form  3'145  grams  of  iron  sulphide. 

A  magnet  will  remove  the  iron  from  the  mixture  of  iron  and 
sulphur  before  it  is  heated ;  but  no  such  separation  is  possible 
after  the  reaction  has  taken  place.  The  definite  compound  iron 
sulphide  is  formed. 

It  should  be  borne  in  mind  that  the  metal  and  sulphur  can 
be  mixed  together  in  any  proportions,  and  easily  separated 
again ;  but  neither  variation  in  proportions  nor  easy  separation 
is  possible  when  the  elements  are  united  in  the  sulphide.  Thus 
a  clear  distinction  can  be  drawn  between  a  mixture  of  two 
elements,  and  a  compound  of  the  same.  Read  the  statement  of 
the  law  on  p.  47. 

Taking  the  atomic  weights  Fe  =  56  and  S  =  32,  and  using 


60  METALLUKGICAL  CHEMISTEY 

the  principles  employed  in  the  last  chapter,  the  formula  for 
iron  sulphide  is  readily  found  with  the  data  supplied  by  the 
last  experiment. 

If  the  ratio  is  not  evident  from  the  numbers  obtained  by 
the  division  with  the  atomic  weights,  the  simple  plan  is  to 
divide  each  of  these  numbers  by  the  lowest  of  them.  This 
always  gives  1  as  the  lowest  member  of  the  ratio  ;  comparison 
is  then  easy. 

The  chemical  change  may  also  be  expressed  by  the 
equation  : 

Fe   +   S   =  FeS 

56       32         88 

which  is  proved  by  the  results  of  Exp.  27. 

For  2  grams  of  iron  require  M45  grams  of  sulphur 
•*•  1          ?)         j         »>  )>  » 


.-.  56        „        „        „  =  32-032  grams  of 

sulphur. 

A  similar  calculation  made  with  the  results  of  Exp.  26  proves 
that— 

2Cu   +   S   =   Cu.2S 
126        32         158 

These  examples  furnish  further  evidence  in  favour  of  the  broad 
statement  that  atomic  weights  of  elements  expressed  in  grams, 
or  any  other  weight  denomination,  contain  the  same  number 
of  atoms.  They  should  be  referred  to  when  difficulties  arise 
as  to  the  meaning  of  an  equation. 

Lead  and  Tin.  —  Similar  experiments  may  be  made  with 
finely  divided  lead  and  tin,  produced  by  vigorously  shaking 
the  molten  metals  in  a  wooden  box  just  as  they  are  about  to 
solidify.  But  in  this  case  some  of  the  vapour  condenses  on 
the  sides  of  the  tube  along  with  the  excess  of  sulphur,  so 
that  a  good  quantitative  result  cannot  be  obtained.  The 


METALS  AND  SULPHUE  61 

compounds  prepared  as  above  are  usually  described  as  artificial 
sulphides,  to  distinguish  them  from  the  naturally  occurring 
ones. 

Natural  Sulphides. — The  majority  of  the  common  metals 
are  found  in  the  earth's  crust  as  sulphides  mixed  with  earthy 
matter,  and  when  the  metals  they  contain  can  be  profitably 
extracted  from  them  they  form  the  useful  ores  of  these  metals. 

Thus,  Galena  contains  lead  sulphide,  PbS.    Copper  Pyrites 

contains  copper  and  iron  sulphides,  Cu2S,  Fe2S3,  and  is  used 
as  an  ore  of  copper.  Iron  Pyrites  is  almost  entirely  a 
sulphide  of  iron,  FeS2.  These  bodies  ofcen  contain  gold  and 
silver. 

ACTION  OF  HEAT  AND  AIR  ON  SULPHIDES. 

Most  sulphides  melt  when  strongly  heated  out  of  contact 
with  air,  and  some  of  them  are  converted  into  vapour;  but  no 
other  change  is  observed.  If  air  is  present,  another  change 
takes  place  where  the  air  comes  into  contact  with  the  sulphide. 
It  is  a  surface  change,  and  proceeds  most  rapidly  when  the 
material  is  finely  divided,  and  the  temperature  is  kept  below 
the  point  at  which  the  particles  soften  and  clot  together.  If 
the  action  is  continued  until  no  further  change  takes  place  the 
sulphur  disappears,  and  its  place  is  largely  taken  up  by 
oxygen.  The  sulphide  is  thus  converted  into  oxide.  There 
are  some  exceptions  to  this  statement,  but  it  is  generally  true. 
It  is  noticed  that  as  the  action  proceeds  the  material  gets 
more  and  more  infusible,  till  finally  it  can  be  raised  to  a  bright 
red  heat  without  partial  fusion  or  clotting.  This  is  because 
oxides  are  generally  more  infusible  than  the  corresponding 
sulphides  from  which  they  are  formed.  The  above  operation 
is  technically  known  as  roasting1.  This  kind  of  work  requires 
a  muffle  furnace  to  be  properly  carried  out,  and  is  fully  dealt 
with  in  works  on  metallurgy;  but  some  useful  information 
can  be  obtained  by  simple  experiments  about  general  and 
intermediate  changes. 


62 


METALLURGICAL  CHEMISTRY 


EXP.  28. — Put  a  little  finely-powdered  iron  pyrites  into  a  small 
test-tube,  and  carefully  heat  the  tube  in  theBunsen  flame.  Sulphur 
vapour  is  given  off,  which  condenses  to  liquid  on  the  sides  of  the 
tube,  and  finally  solidifies.  A  residue  is  left  at  the  bottom  of  the 
tube,  which  resembles  the  iron  sulphide  of  Exp.  27. 

EXP.  29. — Put  a  little  pyrites  into  the  middle  of  a  piece  of  glass 
tube  open  at  both  ends  ;  hold  the  tube  in  a  slanting  position  over 
the  Bunsen  flame,  so  that  the  pyrites  is  heated.  A  pale  blue  flame 
is  seen  over  the  heated  powder.  The  position  of  the  tube  encourages 
air  to  pass  in  at  the  bottom,  through  the  tube,  and  out  at  the  top. 
The  sulphur  driven  out  of  the  pyrites  by  the  heat  burns  in  the 
oxygen  of  the  air  passing  through  the  tube,  and  the  smell  of  burning 
sulphur  is  readily  detected  in  the  gas  issuing  from  the  top. 

It  should  be  noted  here  that  iron  pyrites,  FeS2,  contains 
twice  as  much  sulphur  as  common  iron  sulphide,  FeS,  but 
that  half  of  it  is  more  loosely  held,  and  is  driven  off  when 
the  compound  is  heated,  leaving  the  simpler  compound  as  a 
residue. 


FIG.  16. 

EXP.  30. — Fit  up  the  apparatus  shown  in  Fig.  16.  A  is  a  piece 
of  combustion-tube  about  5  inches  long,  drawn  off  at  one  end,  and 
connected  with  the  bottle  B,  which  contains  a  little  water.  The 
aspirator  C,  which  is  simply  a  large  bottle  with  a  tap  at  the  bottom, 
is  nearly  filled  with  water,  and  connected  with  B  by  a  flexible  tube. 
When  the  tap  at  the  bottom  of  C  is  opened  water  runs  out,  and  air 
passes  through  A  and  B  into  C  to  take  its  place.  The  rate  at  which 
air  flows  through  A  can  be  regulated  by  the  tap,  and  a  steady 
current  thus  kept  up.  If  now  some  iron  pyrites  is  put  into  A, 


METALS  AND  SULPHUR  63 

and  heated  by  a  Bunsen  flame  underneath,  while  a  current  of  air 
is  aspirated  over  it,  a  blue  flame  is  seen  in  the  tube,  and  white  fumes 
appear  in  B.  The  action  may  be  continued  for  ten  minutes,  then 
stopped,  the  bottle  B  removed,  and  the  fumes  blown  out.  The 
water  smells  strongly  of  sulphur  dioxide,  which  proves  that  one  of 
the  changes  resembles  the  burning  of  sulphur  in  air. 

If  the  roasting  is  continued  long  enough  the  whole  of  the 
sulphur  is  practically  removed,  and  the  residue,  which  is 
bluish  black  in  colour,  is  found  to  be  an  oxide  of  iron,  having 
the  formula  Fe203.  Now  part  of  the  sulphur  of  the  pyrites, 
FeS2,  is  driven  off  by  heat  alone,  leaving  FeS.  See  Exp.  28. 
The  remainder  is  removed  by  the  direct  action  of  the  oxygen 
of  the  air.  The  various  changes  are  expressed  by  the  equa- 
tions : 

1.  2FeS2     =    2FeS   +      2S 

Iron  pyrites.     Iron  sulphide.    Sulphur. 

2.  2S  +  202  =  2S02 

3.  2FeS  +  70=  Fe^O.  +  2S02. 

Ferric  oxide. 

Sulphuric  Acid,  H2S04. — By  special  means  it  is  possible 
to  make"  the  sulphur  dioxide  take  up  a  further  proportion  of 
oxygen  from  the  air.  It  is  thus  converted  into  another 
definite  compound,  sulphur  trioxide,  S03,  a  white  volatile  solid, 
which  unites  readily  with  water. 

EXP.  31. — Fit  up  apparatus  similar  to  that  shown  in  Fig.  16,  but 
insert  a  loosely-fitting  plug  of  asbestos,  the  fibres  of  which  have 
been  coated  with  finely- 
divided  platinum.  The 
position  of  the  plug  is 
shown  in  Fig.  17,  Put 
some  coarsely  -  powdered 
iron  pyrites,  or  sulphur,  in 
the  open  end  of  A  (Fig.  16), 
heat  it  with  the  Bunsen 
flame,  and  heat  the  asbes- 
tos at  the  same  time  with 

a  separate  burner.     Aspi-  FIG.  17. 

rate   a   moderate   current 

of  air  through  the  apparatus.  The  sulphur  dioxide  formed  by  the 
roasting  of  the  pyrites,  or  the  burning  of  the  sulphur,  is  drawn, 


64  METALLURGICAL  CHEMISTRY 

together  with  the  excess  of  air,  through  the  hot  platinized  asbestos. 
A  should  not  be  less  than  1  inch  internal  diameter,  and  the  quantity 
of  sulphide  limited.  A  little  added  from  time  to  time  gives  the  best 
results.  As  the  air  is  only  partially  exhausted  when  it  reaches  the 
asbestos,  the  sulphur  dioxide  takes  up  more  oxygen,  and  is  con- 
verted into  the  trioxide.  Dense  white  fumes  now  appear  in  B,  and 
readily  dissolve  into  the  water  present.  The  last  change  is  expressed 
thus : 

SO3  +  H2O=:H2S04. 

Remove  the  bottle  B,  move  it  about  so  as  to  make  the  liquid  run 
over  the  whole  surface,  drain  it  into  a  porcelain  evaporating-dish. 
and  heat  it  carefully  over  the  Bunsen  flame  until  white  fumes 
appear.  When  cold  the  dish  is  found  to  contain  a  small  quantity 
of  an  oily  liquid  with  very  powerful  acid  properties,  and  resembling 
in  every  way  oil  of  vitriol,  or  sulphuric  acid,  H2S04. 

The  pure  acid  is  a  heavy  oily  liquid  of  definite  composition, 
which  mixes  with  water  in  all  proportions.  The  commercial 
acid  usually  contains  a  little  water.  The  manufacture  of  the 
commercial  acid  is  based  on  the  changes  described  above,  but 
oxygen  is  carried  from  the  air  to  the  sulphur  dioxide  by  means 
of  oxides  of  nitrogen  (Chap.  VII.).  The  plant  required  for 
its  economical  production  is  very  extensive,  and  a  detailed 
account  of  the  process  will  be  found  in  any  of  the  larger  works 
on  chemistry. 

It  has  been  assumed  that  in  roasting  a  sulphide  the  whole 
of  the  sulphur  is  converted  into  sulphur  dioxide,  and  its  place 
in  the  sulphide  taken  up  by  oxygen.  In  this  case  an  oxide  of 
the  metal  is  formed  ;  but  if  the  temperature  is  kept  down, 
oxygen  is  actually  absorbed  into  the  mass,  and  a  more  complex 
compound,  containing  oxygen  in  addition  to  the  metal  and 
sulphur,  is  formed. 

EXP.  32. — Put  about  5  grams  of  powdered  copper  sulphide, 
Cu2S,  into  a  fire-clay  roasting-dish,  and  place  the  dish  over  the 
Bunseu  flame.  Stir  the  powder  constantly  with  an  iron  scraper, 
and  regulate  the  flame  so  that  the  mass  just  softens,  but  does  not 
clot  sufficiently  to  stick  to  the  dish.  Continue  the  roasting  for 
twenty  minutes,  or  longer,  if  time  permits.  Transfer  the  residue 
to  a  porcelain  mortar,  and  grind  it  to  a  powder  again.  Boil  the 
powder  for  a  minute  or  two  in  20  c.c.  of  water  contained  in  a  small 


METALS  AND  SULPHUK  65 

beaker.  Filter  the  solution  into  a  clean  porcelain  dish,  and,  if  it  is 
deep  blue  in  colour,  set  it  aside  to  crystallize  ;  if  not,  drive  off  some 
of  the  water  by  boiling  the  solution  over  a  small  flame  before  setting 
it  aside.  Boil  up  a  little  of  the  raw  sulphide  with  water  in  a  test- 
tube,  and  notice  that  the  solution  remains  colourless. 

Blue  crystals  are  obtained,  which  are  readily  seen  to  be 
similar  to  the  crystals  of  commercial  blue  vitriol.  The 
changes  by  which  the  new  compound  is  formed  are  somewhat 
complex,  and  cannot  be  considered  at  this  stage  ;  but  the 
final  result  is  equivalent  to  an  absorption  of  oxygen  by  which 
copper  sulphate,  CuS04,  is  produced. 

Percentage  Composition.—  The  usual  way  of  expressing 
the  composition  of  a  compound,  after  it  has  been  experiment- 
ally obtained,  is  to  put  it  into  the  form  of  percentages.  This 
is  known  as  percentage  composition. 

EXAMPLE.—  To  find  the  percentage  composition  of  copper  sulphide 
from  the  data  given  with  Exp.  26. 

2*507  grams  of  copper  sulphide  contain  2  grams  of  copper  ; 

o 
.*.  1  gram  of  copper  sulphide  contains  nT     £rams  °^  c°PPer  5 


Therefore  the  percentage  composition  is  | 


.'.100  grams         „  „         contain    0          =  79*77    grams    of 

2'&u'        copper. 

Similarly  for  sulphur  —         —  =20*22  grams  of  sulphur. 

79*77 
=  20*22 
99*99 

To  find  the  formula  of  the  compound,  the  percentage  of 
each  element  is  divided  by  its  atomic  weight,  and  each 
quotient  so  obta:ned  is  divided  by  the  lowest  quotient.  The 
second  division  gives  a  series  of  numbers,  from  which  it  is 
easy  to  derive  the  formula.  Thus  — 

7|ir  =  l  '25j  and  ^  ~-~~  °'631  ;  then  o^  =  2  very  nearly'  and 

0*631 


0*631 
The  required  formula  is  Cu2S. 


__  i 


66  METALLUKGICAL  CHEMISTRY 

The  method  of  finding  the  formula  of  a  compound  depends 
upon  :  (1)  the  invariable  proportions  of  the  elements  in  the 
compound ;  (2)  the  existence  of  atoms,  which  for  the  same 
element  are  of  invariable  weight.  If,  then,  it  is  found  that 
certain  weights  of  two  or  more  elements  combine  to  form  a 
compound,  and  the  idea  of  atoms  is  applied,  the  only  inference 
which  can  be  drawn  is,  that  the  weight  of  any  one  of  the 
elements  is  made  up  of  an  enormous  number  of  similar  atoms. 
And  it  is  evident  that  if  the  atomic  weight  is  known,  the 
relation  between  it  and  the  known  weight  of  the  element  in 
the  compound  is  proportional  to  the  number  of  atoms  in  that 
weight. 

79*77 
Thus  the  ratio  =  1  -25  is  proportional  to  the  number  of 

DO 

20-9O 
atoms  in  79-77  parts  of  copper,  and  the  ratio  — ^   =0'631  is 

proportional  to  the  number  of  atoms  in  20'22  parts  of 
sulphur. 

Results  obtained  in  this  way  lead  to  what  are  called 
empirical  formulae,  for  they  only  give  the  ratio  between  the 
numbers  of  the  different  atoms  in  the  molecule  of  the  com- 
pound. Thus  Cu2S  might  be  written  Cu4S2,  Cu6S3,  etc.  It 
is,  however,  the  only  method  of  obtaining  formulae  which  is 
applicable  to  all  compounds  without  exception.  Other  facts 
have  to  be  taken  into  consideration  in  deciding  as  to  the 
actual  number  of  atoms  in  the  molecule  of  a  compound  ;  and, 
as  they  are  not  applicable  generally,  there  are  many  com- 
pounds the  actual  molecular  formula  of  which  are  not 
known. 

SUMMARY. 

The  matter  of  this  chapter  should  receive  careful  attention, 
especially  that  portion  which  refers  to  the  determination  of 
formulae  and  the  formation  of  equations.  So  much  depends 
upon  getting  clear  ideas  of  the  quantitative  relations  between 


METALS  AND  SULPHUR  67 

the  elements  in  a  compound,  that  it  is  hardly  possible  to 
devote  too  much  attention  to  it  at  first.  Some  difficulty  is 
always  experienced  by  those  who  are  not  acquainted  with  the 
use  of  symbols  in  general,  in  their  endeavour  to  grasp  the 
meaning  of  a  formula  or  of  an  equation.  The  facts  of  a 
particular  reaction  and  the  attempt  to  explain  them  are 
sometimes  confused ;  so  that  it  is  a  good  plan  to  get  the  facts 
well  digested  before  attempting  their  explanation. 

The  formation  of  sulphides,  and  the  method  of  finding  their 
formulae,  should  be  carefully  studied.  The  exact  composition 
of  a  compound  can  be  found,  either  by  building  it  up  from  its 
elements  (synthesis),  or  by  separating  it  into  its  elements 
(analysis).  This  is  the  first  step,  and  whichever  method  is 
adopted,  both  the  nature  of  the  elements  in  the  compound 
and  their  proportions  by  weight  must  be  determined.  Very 
often  it  is  possible  to  employ  both  methods  for  the  same 
compound,  in  which  case  they  must  confirm  each  other.  Any 
number  of  experiments,  however  diverse  they  may  be,  if  they 
are  directed  to  the  determination  of  the  composition  of  a 
given  compound,  should  lead  to  the  same  result. 

The  distinction  between  the  facts  of  a  particular  reaction 
and  the  theoretical  considerations  used  in  their  explanation 
is  well  brought  out  by  studying  the  equation  : 

Fe     +     S     =     FeS 
56  32  88 

The  fact  is  that  56  grams  of  iron  will  combine  with  32  grams 
of  sulphur  to  form  88  grains  of  iron  sulphide.  The  theory 
states,  in  scarcely  less  definite  language,  that  for  every  atom 
of  iron  in  56  grams  of  the  metal  there  is  a  companion  atom 
of  sulphur  in  32  grams  of  that  element ;  that  these  atoms  run 
together  in  pairs  to  form  molecules  of  iron  sulphide ;  and 
that  whatever  takes  place  between  a  single  pair  of  atoms  is 
simply  repeated  with  every  other  pair  in  the  reacting 
mass. 

This  is   equivalent   to  saying  that  whatever  may  be  the 

5—2 


68  METALLURGICAL  CHEMISTRY 

number  of  atoms  in  56  grams  of  iron,  there   is  the   same 
number  of  atoms  in  32  grams  of  sulphur. 

In  practice,  then,  the  symbols  represent  measurable  pro- 
portional weights  of  the  elements.  Theoretically  they  repre- 
sent single  atoms,  and  one  side  of  an  equation  shows  a  group 
containing  the  minimum  number  of  atoms  which  can  take  part 
in  the  chemical  change  represented  by  the  equation,  the  other 
the  result  of  the  change.  Then,  however  great  the  reacting 
mass  may  be,  it  is  divided  up  into  similar  groups  as  the  change 
proceeds,  and  any  excess  of  one  kind  of  atoms  over  the  others 
takes  no  part  in  it. 

QUESTIONS. 

1.  What  changes  take  place  when  sulphur  is  heated  out  of 
contact  with  air  ? 

2.  How    may    copper    be    converted    into    its    sulphide  ? 
Describe  the  experiment. 

3.  Write  down  a  chemical  equation,  and  say  in  your  own 
words  what  you  think  it  means. 

4.  How  may  sulphuric  acid   be   prepared  from   sulphur1? 
State  the  general  properties  of  the  acid. 

5.  What  takes  place  when  a  sulphide  is  heated  in  air  ? 

6.  What  is  understood  by  the  percentage  composition  of  a 
compound  1    From  the  example  given  with  Exp.  27  find  the 
weights  of  iron  and  sulphur  in  100  grams  of  the  sulphide,  and 
from  the  percentages  so  obtained  deduce  the  formula  of  the 
compound. 

7.  The  empirical  formula  of  a  compound  does  not  always 
represent  its  molecule.     Why  is  this  1 


CHAPTER  VI 
COMMON  ELEMENTS  AND  COMPOUNDS 

Common  Salt. — This  well-known  substance  is  found  in  the 
earth  in  various  parts  of  the  world.  The  most  extensive 
deposits  in  this  country  are  found  in  Cheshire,  Worcester- 
shire, and  Yorkshire.  It  also  forms  the  bulk  of  the  dissolved 
matter  in  sea  water.  It  is  a  white  solid  which  dissolves 
readily  in  water,  and  crystallizes  from  its  solutions  in  cubes. 
The  solid,  or  its  solution,  has  a  characteristic  saline  taste, 
by  which  it  is  readily  recognised.  As  it  is  easily  purified, 
the  commercial  salt  is  sufficiently  pure  for  most  purposes. 
Common  salt,  or  sodium  chloride,  contains  the  metal  sodium 
and  the  gas  chlorine  in  the  proportion  of  23  parts  of  the  metal 
to  35-5  parts  of  the  gas  by  weight.  Its  formula  is,  therefore, 
NaCl. 

Hydrochloric  Acid. — This  very  important  compound,  as 
prepared  for  use,  is  a  strongly-fuming  liquid  with  very  char- 
acteristic properties.  It  is  known  also  as  muriatic  acid  and 
smoking  salts.  The  pure  compound  contains  1  part  of 
hydrogen  to  35 -5  parts  of  chlorine  by  weight,  and  its  formula 
is  HC1.  It  is  prepared  from  common  salt. 

EXP.  33. — Dry  some  common  salt  by  heating  it  in  an  iron  pan 
over  the  Bunsen  flame.  Make  a  dilute  solution  of  sulphuric  acid 
by  pouring  some  of  the  strong  acid  slowly  into  its  own  volume  of 
water.  Notice  that  the  solution  becomes  very  hot,  and  remember 
that  the  acid  must  always  be  poured  into  the  water,  and  not  the 
water  into  the  acid.  Neglect  of  this  rule  may  cause  an  accident. 
Put  a  little  dry  salt  into  a  dry  test-tube,  add  a  little  of  the  diluted 
acid,  and  warm  the  mixture  over  the  Bunsen  flame.  A  gas  is 


70 


METALLURGICAL  CHEMISTRY 


given  off  which  fumes  strongly  in  the  air,  and  has  a  strong  acid 
taste  and  smell.  A  strip  of  blue  litmus-paper  held  in  the  fumes 
is  turned  red.  The  gas  is  hydrochloric  acid. 

EXP.  34. — Fit  up  the  apparatus  shown  in  Fig.  18.  A  is  a  12-ounce 
flask  fitted  with  a  bung  and  angle-tube.  B  is  a  drying  bottle  con- 
taining a  little  strong  sulphuric  acid,  and  standing  in  a  beaker  of 
water  to  keep  it  cool  C  is  a  gas  cylinder  with  a  delivery-tube.  A 
and  C  are  connected  with  B  by  flexible  rubber  tubes.  D  is  a  long- 
necked  flask  with  a  deli  very- tube,  and  filled  up  to  the  mark  with 
water.  Put  about  10  grams  of  dry  salt  into  A,  and  add  50  c.c.  of 
sulphuric  acid  (1  to  1) ;  replace  the  bung,  and  heat  the  flask  gently 
on  a  piece  of  gauze  over  the  Bunsen  flame.  Gas  is  given  off,  passes 


FIG.  18. 

through  the  strong  acid  in  B,  which  dries  it,  and  collects  in  C  by 
gradually  displacing  the  air.  White  fumes  will  soon  appear  at  the 
mouth  of  the  jar,  but  a  considerable  quantity  of  the  gas  must  escape 
before  the  whole  of  the  air  is  expelled.  When  the  jar  is  thought  to 
be  full  remove  the  delivery-tube,  and  put  a  greased  plate  over  the 
mouth  of  the  jar.  Fill  another  cylinder  with  the  gas,  and  then 
connect  B  with  the  deli  very- tube  in  the  bottle  D,  which  contains 
water  up  to  the  mark  when  the  delivery-tube  is  in  position.  Allow 
the  gas  to  pass  into  D  until  the  water  shows  signs  of  running  back. 
Disconnect  from  B,  and  note  that  the  level  of  the  water  in  D  is  con- 
siderably above  the  mark,  the  delivery-tube  still  being  in  position; 
also  that  the  bottle  feels  warm  to  the  hand. 

The  chemical  change  by  which  the  gas  is  liberated  is  ex- 
pressed by  the  equation : 

NaCl  +  H2S04  =  NaHS04  +  HC1. 


COMMON  ELEMENTS  AND  COMPOUNDS  71 

The  sodium  hydrogen  sulphate  formed  in  the  change  remains 
in  the  flask,  and  is  easily  obtained  as  a  crystalline  solid  from 
the  residue. 

Properties  of  Hydrochloric  Acid. — The  dry  gas  col- 
lected in  the  cylinders  is  colourless  and  transparent.  On 
removing  the  plate  from  a  jar  of  the  gas,  fumes  appear  at  the 
mouth.  The  escaping  gas,  on  coming  into  contact  with  mois- 
ture in  the  air,  dissolves  in  it,  and  causes  its  condensation  as 
an  acid  cloud. 

A  burning  candle  held  in  a  deflagrating  spoon  and  carefully 
lowered  into  the  gas  ceases  to  burn,  but  just  before  the  flame 
disappears  the  blue  portion  of  it  turns  green.  The  gas  is  non- 
combustible,  and  does  not  support  combustion. 

A  little  blue  litmus  solution,  when  poured  into  the  jar  of 
gas  used  in  the  last  experiment,  is  turned  red,  showing  that 
the  gas  has  acid  properties. 

When  a  jar  of  the  gas  is  inverted  over  water  and  the  plate 
removed,  the  water  rushes  up  and  fills  the  jar.  The  gas  is 
very  soluble  in  water,  and  as  it  is  dissolved  the  water  rises 
in  the  jar  to  take  its  place.  The  dissolution  of  the  gas  in  the 
bottle  D  in  Exp.  34  brings  out  two  important  facts  :  (1)  that 
the  volume  of  the  solution  is  greater  than  the  original  volume 
of  water ;  (2)  that  heat  is  developed  during  the  dissolution  of 
the  gas.  To  be  quite  sure  of  the  latter  fact  it  is  necessary  to 
put  the  drying  bottle  B  into  a  vessel  containing  cold  water  to 
ensure  that  cold  gas  passes  into  D.  Also  the  weight  of  the 
solution  is  greater  than  that  of  water  bulk  for  bulk,  and 
increases  with  the  quantity  of  gas  absorbed,  up  to  the  point 
of  saturation.  The  strong  solution  at  ordinary  temperatures 
has  a  density  of  1*2,  and  contains  40  per  cent,  of  HC1.  To 
produce  this  the  water  dissolves  more  than  400  times  its  own 
volume  of  the  gas.  Thus  the  strong  commercial  acid  contains, 
roughly,  40  parts  by  weight  of  the  gas  dissolved  in  60  parts 
by  weight  of  water. 

Large  quantities  of  the  acid  are  obtained  in  the  manufac- 


72  METALLUEGICAL  CHEMISTEY 

ture  of  salt  cake  from  common  salt  and  sulphuric  acid.  The 
operation  is  conducted  in  a  furnace,  and  at  a  higher  tempera- 
ture than  that  which  can  be  used  with  a  glass  flask.  In  this 
case  double  the  quantity  of  salt  is  decomposed  by  the  same 
quantity  of  acid.  The  reaction  takes  place  in  two  stages  : 

1.  NaCl  -f  H2S04      =  NaHS04  +  HCl. 

2.  NaCl  +  NaHS04  =  Na2S04  +  HC1. 

Sodium  hydrogen       Sodium 
sulphate.  sulphate. 

The  gas  is  absorbed  while  passing  through  towers  filled 
with  coke  over  which  water  is  allowed  to  trickle.  The  solu- 
tion is  collected  in  receivers  at  the  bottom  of  the  towers.  The 
crude  acid  is  often  impure,  and  must  be  purified  for  special 
purposes,  but  the  common  acid  is  useful  for  many  operations 
in  which  purity  is  not  essential. 

CHLORINE. 

This  element  is  a  yellowish -green  gas  which  is  present  in 
a  number  of  compounds,  of  which  common  salt  is  the  type. 
It  can  be  obtained  either  from  common  salt  or  from  hydro- 
chloric acid.  Black  oxide  of  manganese,  which  is  used  in  its 
preparation,  is  found  in  the  earth  as  manganese  ore  or  pyro- 
lusite.  Its  formula  is  Mn02.  It  is  a  black,  easily-powdered 
solid. 

EXP.  35. — Well  mix  a  little  dry  salt  with  some  powdered  black 
oxide  of  manganese.  Put  the  mixture  into  a  test-tube,  add  a  little 
dilute  sulphuric  acid,  and  gently  warm  the  tube.  A  yellowish- 
green  gas  is  given  off,  which  has  a  very  irritating  odour.  The  effect 
of  the  small  quantity  which  gets  into  the  air  is  quite  perceptible.  A 
piece  of  moistened  litmus-paper  held  in  the  gas  is  quite  decolourized. 
This  yellowish-green  gas  which  is  liberated  from  the  salt  by  the 
combined  action  of  the  oxide  and  the  acid  is  chlorine,  C12. 

EXP.  36.— Use  the  apparatus  shown  in  Fig.  18.  The  wash-bottle 
need  not  be  kept  cool.  Put  about  20  grams  of  black  oxide  of  man- 
ganese into  A,  add  100  c.c.  of  strong  hydrochloric  acid,  and  heat  the 
flask  gently  on  the  gauze.  Allow  the  gas  to  collect  in  C  until  the 
jar  seems  full  of  a  yellowish -green  body.  Cover  the  mouth  with  a 
greased  plate.  Collect  several  jars  of  the  gas  in  the  same  way. 


COMMON  ELEMENTS  AND  COMPOUNDS  73 

The  experiment  must  be  conducted  in  a  fume-chamber,  but 
if  one  is  not  available  the  gas  should  be  collected  over  water. 
In  that  case  a  delivery-tube  is  passed  directly  from  the  flask  A 
to  a  pneumatic  trough,  and  under  the  mouth  of  an  inverted 
cylinder  filled  with  water.  The  gas  is  somewhat  soluble  in 
water,  but  if  a  rapid  stream  of  it  is  kept  up  no  difficulty  is 
experienced  in  collecting  a  number  of  jars.  Either  the  wet  or 
dry  gas  can  be  used  for  illustrating  most  of  its  properties. 

The  changes  taking  place  in  the  above  experiments  are 
expressed  by  the  equations  : 

1.  Mn02  +  2NaCl  +  3H2S04  =  MnS04+2NaHS04+2H 

Manganese  Manganous 

dioxide.  sulphate. 


Manganous 
chloride. 

Properties  Of  Chlorine.  —  It  is  a  yellowish-green  gas 
when  seen  by  daylight,  or  by  the  light  of  burning  magnesium, 
but  the  colour  is  too  faint  to  be  recognised  by  gaslight.  It 
has  a  very  irritating  odour,  and  care  should  be  taken  not  to 
inhale  it  in  quantity.  Fresh  air  is  the  best  remedy. 

Hydrogen  or  combustible  bodies  containing  hydrogen  burn 
in  this  gas,  and  finely-divided  metals  take  fire  when  brought 
into  contact  with  it. 

Experiments  with  Chlorine.  —  1.  Carefully  lower  a  lighted 
candle,  supported  in  the  bowl  of  a  deflagrating  spoon,  into 
a  jar  of  the  gas.  It  continues  burning,  but  with  a  dull  red, 
smoky  flame,  and  a  thin  deposit  of  soot  is  formed  on  the 
inside  of  the  jar. 

2.  Moisten  a  piece  of  tow  with  turpentine,  hang  it  on 
the  bowl  of  a  deflagrating  spoon,  and  lower  it  into  a  jar 
of  the  gas.  The  turpentine  suddenly  bursts  into  flame, 
and  a  thick  deposit  of  soot  is  formed  inside  the  jar.  The 
candle  and  the  turpentine  contain  carbon  and  hydrogen,  and, 
during  the  combustion  the  hydrogen  only  burns,  the  carbon 
being  set  free  and  deposited  as  soot.  When  the  spoon  is 
lifted  out  of  the  jar  white  fumes  appear  at  the  mouth.  This 


74  METALLURGICAL  CHEMISTEY 

fuming  is  caused  by  the  escape  of  hydrochloric  acid  gas.     The 
change  is  expressed  by  the  equation  : 

H2  +  C12  =  2HC1. 

3.  Grind  some  antimony  to  a  fine  powder  in  a  mortar  ;  put 
it  into  a  small  bottle,  and  tie  a  strip  of  muslin  over  the 
mouth.  Sprinkle  some  of  the  metallic  powder  into  a  jar  of 
chlorine  gas.  A  shower  of  brilliant  sparks  is  seen.  Repeat 
the  last  experiment,  using  finely-divided  copper  in  place  of 
antimony. 

The  changes  are  simple,  and  compounds  resembling  common 
salt  in  general  composition  are  formed.  They  are  called 
chlorides,  and  the  formation  of  copper  chloride  may  be  taken 
as  typical  : 

Cu  +  Cl2  =  CuCl2. 

Cupri<T 
chloride. 

The  conversion  of  metals  into  their  chlorides  is  sometimes 
of  considerable  practical  importance,  and  is  effected  on  the 
large  scale  both  by  free  chlorine  and  by  common  salt. 

Chlorine  is  somewhat  soluble  in  water,  and  forms  a  solution 
which  smells  strongly  of  the  gas.  When  the  solution  stands 
for  some  time  slow  decomposition  of  the  water  takes  place, 
with  formation  of  hydrochloric  acid  and  liberation  of  oxygen  : 


The  change  is  more  rapid  in  daylight,  and  very  rapid  when 
a  body  which  will  take  up  oxygen  readily  is  put  into  the 
solution.  On  this  account  chlorine  is  sometimes  called  an 
oxidizing  agent.  Colouring  matter  of  vegetable  origin  is 
readily  decolourized  by  the  gas  in  the  presence  of  water. 
This  action  is  called  bleaching.  It  is  probable  that  the  oxygen 
liberated  from  the  water  combines  with  the  colouring  matter, 
and  forms  colourless  compounds  with  it.  Mineral  colouring 
matter  is  not  affected. 

4.  Moisten  a  strip  of  turkey-red  cloth  about  half-way  up 


COMMON  ELEMENTS  AND  COMPOUNDS 


75 


and  suspend  it  in  a  jar  of  dry  chlorine.  The  colour  is  rapidly 
discharged  from  the  wet  part,  but  very  slowly  from  the  dry 
portion. 

Chlorine  gas  is  made  on  the  large  scale,  and  absorbed  into 
slaked  lime  to  form  bleaching  powder,  from  which  the  chlorine 
is  readily  liberated  for  use.  Common  salt  and  chlorine  are 
used  in  the  metallurgy  of  copper,  silver,  and  gold. 

Potassium  Chlorate,  KC103.  —This  useful  compound  may 
be  prepared  by  passing  chlorine  gas  into  a  strong  solution  of 
caustic   potash.      A   thistle   funnel, 
bent  as  shown  in  Fig.  19,  is  fitted 
13   to  a  wash-bottle  containing  a  little 
water,  and  dips  into  the  solution  of 
Y  \-  f  "  potash  in  water  (1  to  2).     Chlorine, 

| £  ~]  prepared  as  described  above,  is  passed 

through  the  wash  -  bottle  into  the 
solution,  by  which  it  is  rapidly 
absorbed.  A  crystalline  solid  sepa- 
rates from  the  solution.  When  the  action  is  finished  the 
liquid  is  allowed  to  stand  for  further  separation,  poured  off, 
the  solid  dissolved  in  water,  and  recrystallized  to  purify  it. 

Action  of  Heat. — When  potassium  chlorate  is  heated  it 
splits  up  into  potassium  chloride  and  oxygen,  and  this  reaction 
is  made  use  of  in  the  preparation  of  moderate  quantities  of 
oxygen  for  experimental  purposes. 

EXP.  37. — Put  a  little  powdered  chlorate  into  a  dry  test-tube,  and 
heat  it  strongly  in  the  Bunsen  flame.  It  melts  and  gives  off  bubbles 
of  gas,  which,  when  tested  by  the  introduction  of  a  glowing  splint, 
is  found  to  be  oxygen. 


FIG.  19. 


The  complete  change  is  expressed  thus  : 


2KC10, 


30, 


2KC1     + 

Potassium 
chloride. 

2(39  +  35-5  +  48)  3x32 


245 


96 


76  METALLUKGICAL  CHEMISTEY 

By  mixing  the  salt  with  about  a  quarter  its  weight  of  man- 
ganese dioxide,  and  heating  the  mixture,  the  gas  is  liberated 
much  more  freely  and  at  a  lower  temperature.  The  oxide 
undergoes  very  little  change,  and  its  action  need  not  be  con- 
sidered now. 

EXP.  38. — Thoroughly  dry  some  powdered  potassium  chlorate  in 
a  porcelain  basin  on  the  sand-bath.  Heat  some  powdered  black 
oxide  of  manganese  in  a  porcelain  crucible  over  the  Bunsen  flame. 
Weigh  accurately  2  grams  of  the  dry  chlorate,  and  put  it  into  a  dry 
test-tube  ;  add  about  the  same  weight  of  the  ignited  black  oxide, 
and  well  mix  the  two  by  careful  shaking.  Now  weigh  the  tube  and 
its  contents  accurately,  and  then  heat  it  gently  in  the  Bunsen  flame 
until  no  more  gas  is  liberated.  Two  or  three  minutes'  heating  is 
usually  sufficient.  Allow  the  tube  to  cool  down,  and  reweigh  it. 

EXAMPLE. 

Weight  of  tube  and  mixture  before  heating=15'641  grams 

after         „       =  14'863      „ 


Loss         =00-778  gram 

The  loss  is  due  to  the  escape  of  oxygen  from  the  2  grams  of 
chlorate,  and  may  be  compared  \\ith  the  loss  calculated  by 
using  the  weights  given  in  the  equation  on  p.  75. 

For  245  grams  of  KC1O3  lose  96  grams  of  02 

i  '  ^ 

245 
1         „         „          „     ^-  =  0-784  gram  of  02. 

If  the  test-tube  is  fitted  with  a  rubber  bung  and  delivery- 
tube,  and  is  supported  in  a  horizontal  position  by  a  clip,  the 
evolved  gas  may  be  readily  collected  over  water.  This  method 
is  often  used  to  obtain  oxygen  for  experimental  purposes,  and 
will  be  referred  to  later. 

Bromine  and  Iodine.— These  elements  resemble  chlorine 
in  their  chemical  properties,  and  are  of  considerable  use  in 
the  laboratory. 


COMMON  ELEMENTS  AND  COMPOUNDS  77 

Bromine  is  a  dark  red,  heavy  liquid,  3-18  times  heavier 
than  water.  It  gives  off  a  red  vapour  at  ordinary  temperatures, 
which  is  even  more  irritating  than  chlorine.  It  is  somewhat 
soluble  in  water,  forming  a  red  solution.  It  attacks  water  in 
the  presence  of  bodies  which  will  take  up  oxygen,  and  on  that 
account  is  useful  indirectly  as  an  oxidizing  agent.  The  com- 
monest compound  of  bromine  is  potassium  bromide,  KBr. 

Iodine  is  a  greyish -black  crystalline  solid  which  melts 
readily,  and  gives  off  a  violet  vapour  with  an  irritating  odour. 
It  volatilizes  slowly  at  ordinary  temperatures,  so  that  even  the 
solid  has  an  acrid  smell.  Its  most  important  compound  is 
potassium  iodide,  KI,  which  is  readily  soluble  in  water. 
Iodine  is  only  slightly  soluble  in  water,  but  is  readily  soluble 
in  water  containing  potassium  iodide,  and  is  often  used  in  this 
form.  It  also  dissolves  readily  in  alcohol,  ether,  and  carbon 
bisulphide. 

Compounds  of  bromine  and  iodine  are  present  in  small 
quantities  in  sea-water,  and  are  absorbed  into  sea-plants  during 
their  growth.  When  dry  seaweed  is  heated  to  a  low  red  heat 
in  a  retort,  combustible  gas  is  driven  off,  and  a  black  residue 
left,  which,  on  treatment  with  water,  yields  a  solution  con- 
taining a  considerable  quantity  of  dissolved  matter.  On 
careful  evaporation  in  pans,  potassium  sulphate  crystallizes 
out,  and  is  separated  from  the  hot  solution ;  on  cooling  the  hot 
liquid  a  further  deposit  of  potassium  chloride  is  obtained. 
The  remaining  solution,  called  the  mother  liquor,  contains 
compounds  of  bromine  and  iodine,  and,  after  mixing  with 
sulphuric  acid  and  standing  for  a  time,  is  transferred  to  an 
iron  still,  in  which  it  is  heated  with  black  oxide  of  manganese. 
Iodine  distils  over,  and  is  condensed  in  receivers  in  the  solid 
state.  When  no  more  iodine  passes  over  more  black  oxide  is 
added,  and  the  heating  continued.  Now  bromine  distils  over, 
and  is  collected  in  a  fresh  set  of  receivers  in  the  liquid  state. 
The  bromine  so  obtained  is  usually  sufficiently  pure  for  ordinary 
purposes,  but  the  iodine  needs  purification, 


78  METALLUKGICAL  CHEMISTRY 

It  will  be  noticed  that  the  same  materials  are  used  to 
liberate  bromine  and  iodine  as  are  used  in  Exp.  36  for  the 
liberation  of  the  chlorine  from  common  salt,  and  the  chemical 
changes  taking  place  are  also  similar.  Compare  the  equation 
given  below  with  No.  1  on  p.  73  : 

Mn02  +  2KBr  +  3  H2S04  =  MnS04  +  2K  HS04  +  2H20  +  Br2. 

In  Germany  bromine  is  largely  prepared  by  the  action  of 
chlorine  on  bromides  in  the  mother-liquor  after  the  separation 
of  potassium  chloride  from  the  mixed  potassium  salts  which 
occur  in  the  Stassfurt  salt-beds.  The  reaction  is  simple : 

Cl2+2KBr  =  2KCl  +  Br2. 

Potassium     Potassium 
bromide.        chloride. 

Saltpetre,  also  called  nitre  or  potassium  nitrate,  is  obtained 
in  the  form  of  colourless  crystals,  or  as  a  white  powder.  It  is 
very  soluble  in  water,  and  the  solution  has  a  saline  taste. 
An  efflorescence  which  appears  on  the  surface  of  the  soil  in 
hot  countries,  notably  in  India,  is  found  to  consist  largely  of 
potassium  nitrate.  If  this  surface  layer  is  stripped  off  and 
treated  with  water,  the  nitrate  dissolves,  and  may  be  recovered 
from  the  solution  by  crystallization.  The  formation  of  the 
nitrate  is  brought  about  by  changes  in  the  organic  matter  of 
the  soil  in  the  presence  of  decomposing  rocks  containing 
potassium  compounds.  And  it  is  only  where  these  conditions 
are  present  that  the  compound  is  formed.  There  is  also  a 
method,  based  upon  similar  changes,  for  its  artificial  pro- 
duction. The  corresponding  compound,  Chili  saltpetre  or 
sodium  nitrate,  occurs  as  large  deposits  in  certain  parts  of 
South  America.  This  compound  is  very  similar  in  its  pro- 
perties to  the  potassium  salt,  and,  as  it  is  cheaper,  is  used  for 
some  purposes  in  place  of  the  latter. 

The  formulse  of  the  compounds  are  KN03  and  NaN03 
respectively. 

Nitric  Acid,  HN03. — This  compound  is  a  strongly-fuming 
liquid  with  very  powerful  properties.  It  is  readily  obtained 


COMMON  ELEMENTS  AND  COMPOUNDS 


79 


B 


from  either  of  the  nitrates  described  above  by  acting  upon 
them  with  strong  sulphuric  acid. 

EXP.  39.  —  Fit  up  the  apparatus  shown  in  Fig,  20.  A  is  a  glass- 
stoppered  retort  of  about  8  ounces  capacity.  B  is  a  small  flask 
resting  in  a  bowl  of  water,  and  covered  with  wet  blotting-paper  to 
keep  it  cool. 
Put  about  20 
grains  of  pow- 
dered saltpetre 
into  the  retort, 
and  pour  in 
strong'sulphuric 
acid  through  a 
funnel  until  the 
solid  is  well 
covered.  Ee- 
place  the  stop- 
per, and  gently 
heat  the  mix- 
ture. Eeddish 
vapours  appear 
in  the  retort;  FIG.  20. 

drops  of  liquid 

collect  in  the  stem  and  run  down  into  the  receiver.  Continue  the 
heating  for  twenty  minutes.  Then  remove  the  receiver  and 
examine  the  liquid  which  has  distilled  over. 

The  acid  formed  by  the  reaction  in  the  retort  is  gaseous  at 
the  temperature  of  the  experiment,  and,  passing  in  this  form 
into  the  cooler  stem,  condenses  there,  and  runs  into  the  receiver. 
The  liquid  thus  collected  is  usually  coloured,  which  is  due  to 
the  presence  of  gaseous  impurities  formed  by  the  decomposi- 
tion of  part  of  the  acid  by  heat.  If  a  stream  of  air  is  blown 
through  it  for  some  time,  the  colour  disappears,  and  a  colour- 
less, strongly-fuming  liquid  is  obtained. 

The  change  taking  place  in  the  retort  is  expressed  by  the 
equation  : 


+  H2S04  =  KHS04  +  HNO8 

Potassium 
hydrogen  sulphate. 

39  +  14  +  48  1  +  14  +  48 


101 


63 


80  METALLUEGICAL  CHEMISTEY 

The  potassium  hydrogen  sulphate  is  left  as  a  residue  in  the 
retort.  The  conversion  of  the  nitrate  into  nitric  acid  is  easily 
seen  to  consist  of  an  exchange  of  hydrogen  for  potassium. 
Reference  to  the  table  of  atomic  weights  will  show  that  the 
numbers  under  the  formulae  in  the  equation  represent  the 
molecular  weights  of  the  compounds,  and  also  the  proportion 
between  the  weight  of  nitrate  and  the  weight  of  acid  formed 
from  it.  If  sodium  nitrate  is  used  instead  of  saltpetre,  the 
reaction  is  expressed  thus  : 

NaN03  +  H2S04  =  NaHS04  +  HN03 

Sodium  hydrogen 
sulphate. 

23  +  14  +  48  1  +  14  +  48 


85  63 

On  comparison  it  is  seen  that  85  parts  of  sodium  nitrate 
produce  the  same  weight  of  nitric  acid  as  101  parts  of 
potassium  nitrate.  So  that  if  the  acid  is  the  object  of  the 
preparation,  the  sodium  compound  is  the  more  economical 
to  use. 

In  the  manufacture  of  the  acid  Chili  saltpetre  is  used, 
and  the  reaction  is  conducted  in  iron  retorts  lined  with  fire- 
clay. A  much  higher  temperature  is  employed  and  twice  as 
much  acid  obtained  with  a  given  weight  of  sulphuric  acid. 
This  is  explained  by  the  equation  : 

2NaN08  +  H2S04  =  Na2S04  +  2HN03. 

Sodium 
sulphate. 

The  acid  vapour  is  condensed  in  a  series  of  earthenware 
receivers,  and  afterwards  decolourized  by  blowing  a  current 
of  air  through  the  liquid. 

Properties  Of  Nitric  Acid.— The  ordinary  strong  com- 
mercial acid  has  a  density  of  1  -4,  and  contains  about  65  per 
cent,  of  the  pure  acid,  together  with  35  per  cent,  of  water. 
But  by  carefully  redistilling  this  strong  acid,  mixed  with  half 
its  volume  of  strong  sulphuric  acid,  a  strongly-fuming  liquid 


COMMON  ELEMENTS  AND  COMPOUNDS  81 

is  obtained.  This  is  the  practically  pure  acid,  containing 
100  per  cent.  HN03,  and  having  a  density  of  1*53.  So 
that  pure  nitric  acid  is  a  colourless  liquid,  which  is  readily 
vapourized  and  easily  decomposed.  It  is  much  more  stable 
in  the  presence  of  water,  with  which  it  mixes  in  all  pro- 
portions. 

A  dilute  solution  of  the  acid  turns  blue  litmus  red,  and  has 
the  characteristic  sour  taste  of  acids  in  general.  The  strong 
acid  has  a  very  caustic  action  upon  the  skin,  which  turns 
yellow,  then  black,  and  finally  peels  off.  It  is  a  powerful 
oxidizing  agent,  readily  giving  up  part  of  its  oxygen  to  bodies 
which  will  absorb  that  element.  Sodium  and  potassium 
nitrates  are  similar  to  the  acid  in  this  respect. 

EXP.  40. — Make  a  piece  of  charcoal  red-hot  in  the  Bunsen  flame, 
place  it  on  a  tile,  and  carefully  pour  a  few  drops  of  strong  nitric 
acid  upon  the  incandescent  part.  The  glowing  charcoal  glows  more 
brightly,  showing  that  increased  oxidation  is  taking  place,  due  to 
the  extra  supply  of  oxygen  obtained  from  the  acid. 

EXP.  41. — Melt  a  little  saltpetre  in  a  test-tube  supported  by  a 
clip,  and  drop  in  a  small  piece  of  charcoal.  When  the  charcoal 
begins  to  burn,  remove  the  flame,  and  the  combustion  will  con- 
tinue. The  charcoal  burns  brightly  in  the  oxygen  derived  from  the 
nitrate. 

Repeat  the  last  experiment,  but  drop  into  the  tube  a  small 
piece  of  sulphur  instead  of  charcoal.  The  burner  may  be 
removed  before  dropping  in  the  solid,  as  the  combustion 
commences  at  once  and  continues  vigorously.  An  iron  tray 
should  be  placed  under  the  tube,  as  the  heat  developed  is 
sometimes  sufficient  to  fuse  the  bottom  and  cause  it  to  drop 
out. 

Ammonium  Chloride.— One  of  the  by-products  of  the 
manufacture  of  coal-gas  is  known  as  ammoniacal  liquor.  When 
this  liquid  is  mixed  with  lime  and  heated  in  a  closed  vessel, 
ammonia  gas  is  liberated,  which  is  readily  absorbed  by  strong 
hydrochloric  acid  and  converted  into  ammonium  chloride.  As 
the  acid  liquid  becomes  saturated,  the  chloride  crystallizes  out, 

6 


82  METALLURGICAL  CHEMISTEY 

is  collected,  purified,  and  put  on  the  market  as  sal-ammoniac. 
The  reaction  is  : 

NH3  +  HC1  =  NH4C1. 

Ammonia.  Ammonium 

chloride. 

The  pure  compound  is  a  white,  crystalline,  odourless  solid, 
readily  soluble  in  water.  It  passes  directly  from  the  solid  to 
the  gaseous  state  when  heated.  In  the  gaseous  state  its  mole- 
cules are  largely  split  up  into  ammonia  and  hydrochloric  acid 
molecules  : 

NH4C1  =  NH3  +  HC1. 

But  these  recombine  as  the  temperature  falls,  and  the  original 
solid  is  reproduced : 

NH3  +  HC1  =  NH4C1. 

This  splitting  up  of  the  molecules  at  or  above  a  definite  tem- 
perature, and  their  recombination  below  that  temperature,  is 
known  as  dissociation.  The  last  change  is  readily  shown 
by  filling  two  gas-jars  with  the  dry  gases  and  bringing  them 
together  mouth  to  mouth.  The  white  chloride  forms  rapidly 
and  deposits  on  the  sides  of  the  jars. 

Ammonia  Gas,  NH3. — This  compound  has  a  very  pungent 
odour,  and  its  presence  is  very  readily  detected,  although  it  is 
quite  transparent  and  colourless. 

EXP.  42. —  Mix  together  5  grams  of  ammonium  chloride  and 
4  grams  of  quicklime ;  put  the  mixture  into  a  test-tube  fitted 
with  a  bung  and  a  straight  delivery-tube  about  8  inches  long  ; 
fix  the  test-tube  in  a  clip  in  the  vertical  position,  and  invert  a  wide 
test-tube  over  the  delivery-tube  so  as  to  rest  on  its  end.  Now 
heat  the  mixture,  and  allow  the  gas  to  collect  in  the  inverted  tube. 
When  the  tube  is  full,  which  is  readily  detected  by  the  rapid  escape 
of  the  gas  from  the  mouth,  replace  it  by  a  fresh  one,  and  make  an 
experiment  with  the  full  tube. 

The  heating  of  the  mixture  need  not  be  continued  during 
the  testing  of  the  gas.  Tests  may  be  made  to  prove  the 
following  statements : 

1.  If  a  piece  of  red  litmus-paper  is  held  in  the  gas,  it  is 
turned  blue.  The  gas  has,  therefore,  alkaline  properties. 


COMMON  ELEMENTS  AND  COMPOUNDS  83 

2.  If  a  burning  match  or  taper  is  put  into  an  inverted  tube 
filled  with  the  gas,  the  flame  disappears,  and  the  gas  does  not 
take  fire.     It  is  non-combustible,  and  a  non-supporter  of  com- 
bustion under  ordinary  circumstances.     If  the  gas  is  heated 
strongly,  or  is  mixed  with  oxygen,  it  burns  with  a  somewhat 
feeble,    yellowish    flame;    water    is    formed,   and    nitrogen 
liberated. 

3.  If  a  well-filled  tube  of  the  gas  is  placed  mouth  down- 
wards in  a  dish  of  water,  the  liquid  rises  in  the  tube  to  take 
the  place  of  the  dissolved  gas,  which  is  very  soluble  in  water. 
A  piece  of  red  litmus-paper  dipped  into  the  water  before  the 
dissolution  is  unchanged  ;  but  if  dipped  in  after,  is  turned 
blue. 

4.  If  a  piece  of  clay  tobacco  pipe  is  connected  with  the 
delivery-tube   by   flexible   rubber,  and   its   end   is   strongly 
heated  in  the  Bunsen  flame,  while  a  current  of  gas  is  driven 
through  it  by  heating  the  mixture,  the  gas  burns  at  the  hot 
end  of  the  clay  pipe.     The  change  is  : 


The  solubility  of  ammonia  is  its  most  remarkable  property. 
Cold  water  will  absorb  upwards  of  1,000  times  its  own  volume 
of  the  gas.  The  ordinary  ammonia  liquor  is  simply  a  strong 
solution  of  the  gas,  and  has  many  useful  applications.  The 
alkaline  character  of  the  solution,  and  the  readiness  with 
which  the  gas  escapes  from  it,  has  earned  for  the  compound 
the  name  of  the  volatile  alkali. 

The  principal  fixed  alkalies  are  caustic  potash  and  caustic 
soda,  and  their  solutions  in  water  have  properties  very  similar 
to  those  of  ammonia  solution.  But  they  are  both  solids, 
which  can  be  melted  and  raised  to  a  high  temperature  without 
decomposition.  The  result,  then,  of  evaporating  separate 
solutions  of  these  three  alkalies  would  be  a  clean  vessel  for 
the  first,  and  solid  residues  for  the  other  two.  All  three, 
however,  neutralize  acids,  turn  litmus  blue,  or  methyl  orange 
yellow,  and  have  a  caustic  action  on  the  skin. 

6—2 


84  METALLURGICAL  CHEMISTRY 

The  formulae  for  caustic  soda  and  caustic  potash  are  NaHO 
and  KHO,  and,  by  analogy,  the  formula  of  the  compound  of 
ammonia  supposed  to  be  present  in  its  solution  is  written 
(NH4)HO,  orNH4HO. 

Using  the  systematic  names,  the  following  comparison  can 
be  made : 

NaHO,  sodium  hydroxide. 
KHO,  potassium  hydroxide. 
(NH4)HO,  ammonium  hydroxide. 

The  group  (NH4),  although  it  has  no  free  existence,  plays 
an  important  part  in  ammonium  compounds.  It  resembles  an 
atom  of  sodium  or  potassium  in  its  chemical  functions,  and  is 
called  a  compound  radicle  on  that  account. 

SUMMARY. 

Common  salt,  saltpetre,  and  black  oxide  of  manganese  are 
found  ready  to  hand  in  Nature,  and,  in  conjunction  with 
sulphuric  acid,  are  used  in  the  preparation  of  other  equally 
important  compounds  and  elements.  The  common  acids  are 
exceptionally  important  in  the  treatment  of  metals.  Ammo- 
niacal  gas  liquor  is  the  usual  source  of  ammonia  and  ammonium 
compounds. 

QUESTIONS. 

1.  Explain  what  takes  place  when  common  salt  is  heated 
with  moderately  strong  sulphuric  acid.    Sketch  the  apparatus 
used. 

2.  How  is  chlorine  gas  prepared,  and  what  are  its  most 
important  properties  1     Describe  briefly  the  properties  of  its 
companion  elements. 

3.  What   happens  when   potassium   chlorate   is   heated — 
(a)  alone,  (b)  mixed  with  black  oxide  of  manganese  1 

4.  How  is  nitric  acid  prepared  ? 

5.  Give  a  short  description  of  the  three  common  acids. 

6.  Write  a  short  description  of  ammonia  and  its  compounds. 


CHAPTER  VII 
METALS  AND  ACIDS 

THE  common  forms  of  hydrochloric,  nitric,  and  sulphuric  acids 
are  now  familiar,  and  their  more  evident  properties  have  been 
sufficiently  studied  to  enable  them  to  be  handled  intelligently. 
The  relation  between  metals  and  the  common  acids  is  an 
important  one,  and  the  study  of  the  reactions  which  take 
place  when  they  are  mixed  together  can  now  be  profitably 
undertaken.  In  many  cases  a  very  marked  change  takes 
place,  which  is  readily  noticed,  and  often  easily  investigated. 
But  in  some  cases  no  change  is  effected,  and  the  metal  and 
acid  may  be  left  together  for  any  length  of  time  without 
alteration. 

EXP.  43. — Put  small  pieces  of  zinc,  tin,  and  platinum  into  three 
test-tubes,  and  add  to  each  5  c.c.  of  dilute  nitric  acid.  Warm  the 
tube,  if  necessary,  to  start  the  action,  and  note  'all  the  changes 
which  take  place. 

The  acid  which  is  added  to  the  zinc  boils  up,  and  a  rapid 
escape  of  bubbles  of  gas  is  observed.  The  metal  entirely 
disappears. 

The  tin  changes  to  a  white  powder,  and  a  red  gas  with  an 
acrid  odour  is  given  off. 

The  platinum  is  apparently  unchanged,  and,  if  weighed 
before  and  after  the  experiment,  is  found  to  be  the  same. 

Two  practical  problems  now  present  themselves  :  The  first 
is  to  learn  as  much  as  possible  about  the  liberated  gas  ;  the 
second  is  to  find  out  what  happens  to  the  metal. 

85 


METALLUKGICAL  CHEMISTKY 


To  FIND  THE  VOLUME  OF  GAS  EVOLVED  WHEN  A  WEIGHED 
PIECE  OF  METAL  is  DISSOLVED  IN  AN  EXCESS  OF 
DILUTE  ACID. 

The  apparatus  shown  in  Fig.  21  is  to  be  fitted  up.  A  is  an 
8-ounce  flask  fitted  with  a  rubber  bung  and  an  angle-tube. 
B  is  a  12-ounce  bottle  fitted  with  a  rubber  bung  and  two  angle 
tubes.  C  is  a  gas-cylinder  of  about  500  c.c.  capacity.  B  is 
connected  with  A  and  with  the  delivery-tube  d  by  pieces 
of  sound,  well-fitting  rubber  tube  about  4  inches  long.  These 

flexible  connec- 
tions allow  the 
bung  to  be  re- 
moved from  or 
inserted  in  the 
neck  of  A,  and 
the  delivery -tube 
d  to  be  raised 
above  B  without 
disturbing  the 
apparatus.  Any  form  of  pneumatic  trough  may  be  used,  but 
one  with  a  beehive  shelf  is  very  convenient.  The  apparatus 
must  be  perfectly  gas-tight,  so  that  care  should  be  exercised 
in  the  selection  of  the  fittings  and  in  the  fitting  up,  or  poor 
results  will  be  obtained.  When  ready,  the  apparatus  may  be 
used  for  several  experiments  in  succession,  as  a  single  experi- 
ment requires  a  few  minutes  only  for  completion. 

METALS  AND  HYDROCHLORIC  ACID. 

Zinc — EXP.  44. — Use  the  apparatus  shown  in  Fig.  21.  Add 
a  little  water  to  the  flask  B,  shake  well  so  as  to  wet  the  sides, 
allow  it  to  stand  for  a  minute,  and  then  pour  out  the  excess  of 
water.  Remove  the  bung  from  the  flask  A,  and  add  25  c.c.  of 
dilute  hydrochloric  acid  (1  to  1).  Cut  a  strip  of  zinc  foil  about 
1  gram  in  weight,  and  weigh  it  accurately.  Fill  the  cylinder  C, 
invert  it  on  the  shelf  of  the  trough,  and  arrange  the  end  of  the 
delivery-tube  under  its  mouth.  Examine  the  connections  to  see 
that  they  are  well  made,  as  the  experiment  is  a  failure  if  the  whole 
of  the  expelled  gas  is  not  collected  in  C.  Drop  the  weighed  metal 


FIG.  21. 


METALS  AND  ACIDS  87 

into  A,  and  rapidly  insert  the  bung.  If,  when  the  action  has 
ceased,  the  flask  A  is  warm,  due  to  the  heat  evolved  during  the 
dissolution  of  the  metal  in  the  acid,  place  it  in  a  shallow  pan 
containing  water  which  has  been  standing  in  the  laboratory  for 
some  time.  This  will  quickly  bring  the  flask  and  its  contents 
back  again  to  the  original  temperature,  especially  if  the  flask  is 
moved  about  in  the  water.  When  the  level  of  the  water  in  the 
delivery-tube  is  constant,  raise  the  tube  and  allow  the  liquid  to 
run  back  into  B.  Pour  the  water  from  B  into  a  50  c.c.  measure, 
and  note  its  volume.  This  is  the  measure  of  the  excess  of  gas 
driven  into  C  by  the  increase  in  the  temperature  in  A.  Now 
moisten  a  narrow  strip  of  filter-paper,  and  fix  it  on  C  a  little  above 
the  water-level  in  the  cylinder  ;  depress  the  cylinder  in  the  trough 
until  the  water-level  is  the  same  inside  and  outside  ;  move  the 
paper  till  its  lower  edge  marks  this  level.  Kemove  the  cylinder 
from  the  trough,  and  put  a  lighted  taper  to  its  mouth.  The  ^gas 
is  a  mixture  of  air  and  hydrogen,  and  burns  with  a  slight  explosion. 
Now  fill  the  cylinder  to  the  mark  with  water,  rinse  out  the  500  c.c. 
measure,  pour  in  the  water  from  the  cylinder,  and  note  its  volume. 
This  is  the  measure  of  the  volume  of  gas  collected  in  C,  and,  there- 
fore, of  the  volume  of  gas  liberated  from  the  acid  by  the  weighed 
quantity  of  zinc,  which  must  displace  that  volume  from  A  and  B. 

Further  information  about  the  experiment  is  obtained  by 
noting  the  height  of  the  barometer  column,  which  gives  the 
pressure  of  the  gas,  and  by  taking  the  temperature  of  the 
water  in  the  trough  with  a  thermometer,  which  is  also  the 
same  as  the  temperature  of  the  gas  collected  over  it.  For  the 
volume  of  a  gas  varies  considerably  with  variations  of  its 
temperature  and  pressure,  and  is  increased  by  the  presence  of 
water  vapour  in  it.  In  accurate  determinations  it  is  necessary 
to  reduce  this  volume  to  that  of  the  dry  gas  at  the  normal 
temperature  and  pressure  (N.T.P.).  But  the  beginner  may 
leave  this  for  a  time.  A  series  of  experiments  has  shown  that 
an  approximate  correction  for  the  volume  of  a  gas  collected 
over  water  under  ordinary  laboratory  conditions  is  obtained 
by  multiplying  the  observed  volume  by  0'916. 

EXAMPLE. — Weight  of  zinc  taken  1  gram. 

Volume  of  gas  collected  in  C  =     375  c.c. 
,,          water  in  B  =         3   ,, 

,,          gas  liberated  =     372   ,, 

Then  372  x  0'916  =  340'7  c.c.  =  the  volume  of  dry  hydrogen 
at  N.T J?.  expelled  from  hydrochloric  acid  by  1  gram  of  zinc., 


88  METALLUEGICAL  CHEMISTKY 

The  weight  of  a  litre  of  hydrogen  has  been  determined  with 
considerable  accuracy,  and  may  be  used  to  find  the  weight  of 
any  given  volume  of  the  dry  gas  at  N.T.P. 

1  litre  =  1000  c.c.  =  0-0896  gram. 

Now,  the  volume  of  hydrogen  liberated  by  1  gram  of  zinc 
having  been  found,  it  is  easy  to  calculate  the  volume  liberated 
by  65  grams  of  the  metal—  i.e.,  by  an  atomic  weight  of  zinc  in 
grams. 

340-7x65  =  22145-5  c.c. 

And,  as  each  1,000  c.c.  of  this  gas  weighs  0*0896  gram,  the 
weight  of  the  whole  is  22145-5  x  0-0896  +  1000=  1-984. 

The  number  thus  obtained  from  a  rough  experiment  is 
sufficiently  near  the  whole  number  2  to  suggest  that  an 
atomic  weight  of  zinc  in  grams  liberates  2  atomic  weights 
of  hydrogen  in  grams.  From  the  formula  HC1  it  is  seen  that 
2  molecular  weights  of  hydrochloric  acid  in  grams  are 
required  to  supply  the  2  grams  of  hydrogen,  and  as  the  whole 
of  the  chlorine  combines  with  the  zinc,  the  complete  change  is 
expressed  by  the  equation  : 

Zn  +  2HC1    =   ZnCl2  +  Ho 
65  2(1  +  35-5)     65  +  71         2 


Magnesium  —  EXP.  45.  —  Weigh  0-3  gram  of  magnesium  ribbon 
or  wire.  Add  50  c.c.  of  dilute  hydrochloric  acid  (1  in  20).  Make 
the  magnesium  into  a  coil,  drop  it  into  the  flask,  and  quickly 
replace  the  bung.  The  gas  comes  off  rapidly.  Follow  the  instruc- 
tions given  for  EXP.  44,  and  make  the  necessary  measurements. 
Prove  that  the  liberated  gas  is  hydrogen. 

EXAMPLE.—  Weight  of  magnesium  used  =    0*3  gram. 
Volume  of  gas  in  C  =  312  c.c. 

water  in  B  =       7    ,, 


„          gas  liberated        =305   „ 

NOTE.— The  principles  upon  which  the  above  calculations  are  based  will 
be  found  in  the  Appendix. 


METALS  AND  ACIDS  89 

Then   305x0-916  =  279-4.      Compare   this   result   with    the 
equation  : 

Mg  +  2HCl«MgCla+H2. 

24  22320  c.c. 

Iron.  —  The  dissolution  of  this  metal  is  less  rapid  than  that 
of  zinc  or  magnesium,  and,  when  the  evolution  of  the  gas 
slows  down,  it  is  necessary  to  heat  the  flask.  This  can  be 
done  by  placing  it  on  gauze  over  a  circular  Bunsen  burner. 
The  flexible  connection  between  A  and  B,  Fig.  21,  enables 
this  to  be  readily  done.  The  heating  is  continued  over  a  small 
flame  until  bubbles  of  gas  cease  to  escape  from  the  delivery- 
tube,  or  escape  very  slowly.  As  commercial  iron  contains 
impurities,  the  volume  of  gas  liberated  by  a  given  weight  of 
the  metal  falls  a  little  short  of  that  which  would  be  obtained 
by  using  the  pure  metal. 

EXP.  46.  —  Weigh  0'7  gram  of  clean  iron  filings  ;  transfer  it  to 
the  flask  A,  Fig.  21  ;  add  25  c.c.  of  strong  hydrochloric  acid,  and 
rapidly  replace  the  bung.  When  the  action  slows  down,  heat  the 
flask  gently  until  the  dissolution  is  finished.  Cool  the  flask  and 
make  the  necessary  measurements.  Prove  that  the  gas  is  hydrogen. 
During  the  heating  a  considerable  volume  of  air  will  be  expelled 
from  the  flask,  and  a  much  larger  quantity  of  water  will  run  back 
into  B.  When  the  bung  is  removed  from  A  an  offensive  smell 
will  be  noticed,  which  is  due  to  the  action  of  the  acid  on  the 
combined  carbon  present  in  the  iron. 

EXAMPLE.  —  Weight  of  iron  =     0'7  gram. 

Volume  of  gas  in  C  =     464  c.c. 

„          water  in  B        =     168  „ 

„          gas  liberated     =     296   „ 
Compare  the  result  with  the  equation  : 


Tin.  —  This  metal  dissolves  even  less  rapidly  than  iron. 
The  most  suitable  form  for  general  use  is  the  thin  sheet  metal, 
which  is  readily  obtained  by  rolling  down  a  piece  of  grain  tin, 
or  by  hammering  it  out  on  a  clean  anvil. 


90  METALLUEGICAL  CHEMISTKY 

EXP.  47.—  Weigh  about  1-3  grams  of  tin  ;  add  20  c.c.  of  strong 
hydrochloric  acid  to  A,  Fig.  21  ;  put  in  the  metal,  and  heat  the 
flask  gently  until  the  dissolution  is  complete.  Measure  and 
examine  the  gas  collected  in  C. 

The  volume  of  gas  obtained  is  a  little  greater  than  that 
given  by  the  equation  : 

Sn  +  2HCl=SnCl2  +  H2. 

EXAMPLE.  —  Weight  of  tin  =     1-3  grams. 

Volume  of  gas  in  C  =     390  c.c. 

,,          water  in  B          =     117    ,, 

„          liberated  gas      =     273    „ 

Aluminium.—  The  thin  sheet  metal  is  readily  obtained, 
and  dissolves  in  moderately  strong  hydrochloric  acid  without 
the  application  of  heat.  The  change  is  expressed  by  the 
equation  : 


EXP.  48.—  Weigh  0'3  gram  of  aluminium  ;  add  15  c.c.  of  dilute 
hydrochloric  acid  (1  to  1)  to  A,  Fig.  21,  and  put  in  the  weighed  metal. 
The  action  is  sufficiently  rapid  to  finish  without  heating  the  flask, 
and  very  little  water  runs  back  into  B.  Measure  and  test  the  gas 
collected  in  C. 

EXAMPLE.  —  0'318  gram  of  aluminium  gave  426  c.c.  of  hydrogen  ; 
or  426x0-916=390-2  c.c.  of  the  dry  gas  at  N.T.P. 

Then,  since  1,000  c.c.  of  hydrogen  weigh  0*0896  gram,  the  weight 
of  hydrogen  which  would  be  liberated  by  the  dissolution  of  an 
atomic  weight  of  the  metal  in  grams  is  given  by  : 

390-2x0-0896x27 
'    0-318  x  1000  —2'97  =  3  nearly. 

So  that   1   atomic  weight  of  aluminium  displaces  3  atomic 
weights  of  hydrogen  from  the  acid. 

The  metals  used  in  the  above  experiments  are  those  which 
are  most  readily  acted  upon  by  hydrochloric  acid.  Some 
common  metals  are  dissolved  very  slowly,  and  others,  again, 
are  not  attacked  by  the  acid,  even  on  prolonged  heating.  A 
metal  which  dissolves  slowly  may  have  its  rate  of  dissolution 
increased  by  passing  a  current  of  air  through  the  acid  liquid. 


METALS  AND  ACIDS  91 

This  is  no  doubt  due  to  the  formation  of  oxide  of  the  metal  by 
the  oxidizing  action  of  the  air,  and  its  dissolution  by  the  acid. 

It  may  be  stated  generally  that  hydrogen  gas  is  set  free 
when  a  metal  dissolves  rapidly  in  either  hot  or  cold  hydro- 
chloric acid. 

The  compounds  formed  by  the  replacement  of  hydrogen 
in  the  hydrochloric  acid  molecules  by  the  respective  metals 
are  called  chlorides.  They  are  readily  separated  from  the 
solutions  by  evaporating  the  water  and  excess  of  acid,  or  by 
crystallization.  (See  Exps.  52  and  53.) 

METALS  AND  SULPHURIC  ACID. 

When  hydrochloric  acid  reacts  with  metals  the  general 
change  is  a  simple  one,  corresponding  to  the  simple  character 
of  the  acid  itself.  But  with  sulphuric  acid  there  is  the  possi- 
bility of  secondary  changes,  for,  if  its  molecules  are  regarded 
as  consisting  of  the  two  groups  H2  and  (S04),  the  second  group 
may  undergo  changes  by  which  new  compounds  are  formed. 
The  simplest  kind  of  change  is  when  the  group  (S04)  moves 
into  or  out  of  molecules  in  much  the  same  way  as  an  atom  of 
an  element  does.  This  is  the  simple  explanation  of  the 
ordinary  dissolution  of  a  metal  in  the  dilute  acid  by  which 
hydrogen  is  liberated.  But  when  the  acid  is  hot  and  concen- 
trated, another  reaction  becomes  prominent,  by  which  sulphur 
dioxide  gas  is  set  free.  The  probable  explanation  of  this  is 
that,  at  the  temperature  used,  a  considerable  proportion  of  the 
acid  molecules  are  split  up,  or  dissociated  into  (H2)  and 
(S04)  groups,  and  that  the  hydrogen  set  free  by  the  action  of 
the  metal  on  other  acid  molecules  helps  to  attack  the  (S04) 
groups,  and  abstracts  oxygen  from  them.  This  is  expressed 
in  the  equations  : 


2. 

The  experiments  with  dilute  sulphuric  acid  are  made  in 
exactly  the  same  way  as  with  hydrochloric  acid,  so  that  a 


92  METALLUKGICAL  CHEMISTRY 

detailed  description  will  not  be  required.  The  dilute  solution 
is  made  by  pouring  the  measured  quantity  of  strong  acid  into 
the  measured  quantity  of  water.  The  mixture  is  then  stirred 
with  a  glass  rod,  and  allowed  to  cool.  The  apparatus  shown 
in  Fig.  21  is  to  be  used. 

Zinc—  EXP.  49.—  Use  about  1  gram  of  the  thin  sheet  metal 
and  20  c.c.  of  dilute  sulphuric  acid  (1  in  6).  Measure  and  test  the 
liberated  gas. 

If  the  zinc  is  of  good  quality  the  action  is  very  slow  at  first, 
and  it  may  be  necessary  to  warm  the  flask  a  little. 

EXAMPLE.—  1  gram  of  zinc  gave  377  x  0-916=345-3  c.c.  of 
hydrogen. 

This  result  may  be  compared  with  that  obtained  by  cal- 
culation from  equation  No.  1  on  p.  19. 

Magnesium—  EXP.  50.—  Use  about  0'3  gram  of  the  metal  ribbon 
and  50  c.c.  of  dilute  sulphuric  acid  (1  in  50).  Measure  and  test  the 
liberated  gas. 

EXAMPLE.—  0-3  gram  of  the  metal  gave  305xO'916=280-4  c.c. 
of  gas. 


Iron—  EXP.  51.—  Use  about  0-7  gram  of  the  metal  and  20  c.c. 
of  the  dilute  acid  (1  in  6).  Heat  the  flask  gently,  as  described  in 
Exp.  46. 

EXAMPLE.—  0-7  gram  of  iron  gave  296  x  0'916=27ri  c.c.  of  gas. 
Fe  +  H2S04=FeS04  +  H2. 

On  comparing  the  volume  of  gas  obtained  with  that 
determined  by  calculation  from  the  equation,  it  is  found  to 
be  deficient.  This  is  due  to  the  impurities  in  the  commercial 
metal. 

A  comparison  of  the  results  obtained  above  shows  that  the 
volume  of  gas  liberated  depends  upon  the  weight  of  metal 
used,  and  not  upon  the  nature  of  the  acid. 


ACID. 

ZN  1  GRAM. 

Mo  0-3  GRAM. 

FE  0'7  GRAM. 

HC1    .. 
H2S04 

3407  c.c. 
345-3    „ 

279-4  C.C. 
2804    ,, 

i 

271  '2  c.c. 
271-1    „ 

METALS  AND  ACIDS  93 

The  other  common  metals  are  either  insoluble  or  dissolve 
with  difficulty  in  dilute  sulphuric  acid. 

The  compounds  formed  by  the  replacement  of  the  hydrogen 
in  the  acid  molecules  by  the  metals  are  called  sulphates, 
and  it  will  be  useful  at  this  stage  to  follow  up  the  disappear- 
ing metals  and  learn  something  about  the  bodies  they  pass 
into. 

EXP.  52.  —  Dissolve  2  grams  of  zinc  in  20  c.c.  of  dilute  sulphuric 
acid  (1  to  7)  contained  in  a  small  beaker.  Pour  the  clear  solution 
into  a  porcelain  basin  ;  raise  it  to  boiling  over  the  Bunsen  flame, 
and  set  the  basin  aside  until  the  next  lesson.  Then  drain  away 
the  liquid,  and  dry  the  separated  solid  on  a  piece  of  filter-paper. 

Repeat  the  experiment  with  2  grams  of  clean  iron  filings. 

The  solids  thus  obtained  consist  of  crystals  of  zinc  and 
iron  sulphates,  and  contain  the  remnants  of  the  sulphuric 
acid  molecules,  together  with  water  of  crystallization.  The 
formulae  of  the  salts  are  :  ZnS04,  7H20  and  FeS04,  7H20. 

EXP.  53.  —  Measure  8  c.c.  of  dilute  sulphuric  acid  (1  to  12)  into 
a  large  porcelain  crucible  which  has  been  carefully  weighed  ;  cover 
the  crucible  with  a  watch-glass,  and  add  0'3  gram  of  magnesium,  a 
small  piece  at  a  time.  When  the  whole  is  dissolved,  drain  the 
watch-glass  into  the  crucible,  and  wash  off  any  __ 
adhering  liquid  with  a  fine  jet  of  water  from  a 
wash-bottle.  Place  the  crucible  on  a  pipe-clay 
tripod,  so  that  the  bottom  is  just  clear  of  the 
sand  in  the  sand-bath,  and  allow  the  liquid  to 
evaporate  to  complete  dryness.  Heat  the 
crucible  over  the  Bunsen  flame,  at  first  with  jrIGt  22. 

the    lid   on   in   case   of    spurting,   until   white 
fumes  cease  to  be  given  off.     Allow  the  crucible  to  cool,  and  weigh 
it.     The  final  heating  must  be  done  in  a  fume-chamber. 

EXAMPLE.  —  Weight  of  crucible  +  residue  =   18'665  grams. 

=   17-182       „ 
„        =   1'483 
From  the  equation  : 

Mg   +  H2S04  =        MgS04     +     H2 
24  24  +  32  +  64 


24  grams  of  magnesium  give  120  grams  of  residue.     Therefore 

1  20  x  0'3 
0'3  grain  of  magnesium  should  give  —  z  —  =  1*5  grams. 


94 


METALLURGICAL  CHEMISTRY 


Thus  the  weight  of  sulphate  formed  is  just  as  definite  as  the 
volume  of  hydrogen  liberated.  Similar  experiments  may  be 
made  with  other  metals  and  acids,  with  similar  results. 

PREPARATION  OF  HYDROGEN  GAS. 

The  dissolution  of  zinc  in  dilute  hydrochloric  or  sulphuric 
acid  may  be  used  as  a  source  of  hydrogen  for  experiments 
to  demonstrate  its  properties.  If  the  apparatus  described 
on  p.  24  is  not  available,  this  method  is  to  be  employed  for 
obtaining  a  quantity  of  the  gas. 

EXP.  54. — A  is  a  12-oz.  bottle  fitted  with  a  rubber  bung,  through 

which  passes  a  thistle  funnel,  B, 
and  a  delivery  tube.  About  10 
grams  of  zinc  is  put  into  A, 
covered  with  water,  and  strong 
hydrochloric  acid  added  through 
B,  a  little  at  a  time.  The  gas  is 
thus  liberated  steadily,  and  is 
collected  in  the  cylinder  C.  The 
first  jar  of  gas  contains  some  air 
displaced  from  the  generating- 
bottle  ;  the  gas  collected  after  this 
is  sufficiently  pure  for  ordinary 
experiments. 

Experiments  for  demonstrating  the  properties  of  the  gas  are 
described  on  p.  25. 


B 


PREPARATION  OF  SULPHUR  DIOXIDE. 

As  the  action  of  concentrated  sulphuric  acid  on  metals  is 
not  susceptible  of  easy  quantitative  experiments,  only  one 
example  will  be  given,  and  that  the  one  commonly  used  to 
obtain  sulphur  dioxide  in  sufficient  quantity  to  demonstrate 
its  properties. 

EXP.  55. — Fit  up  the  apparatus  shown  in  Fig.  24.  A  is  a  12-ounce 
conical  flask  ;  B  is  a  wash- bottle  containing  a  little  sulphuric  acid 
(2  parts  of  acid  to  1  part  of  water)  ;  C  is  a  dry  gas-cylinder.  Put 
about  10  grams  of  sheet  copper  into  A,  and  add  25  c.c.  of  strong 


METALS  AND  ACIDS 


95 


FIG.  24. 


sulphuric  acid.     Heat  the  flask  until  the  gas  comes  off  freely,  and 

then  regulate  the  flame  so  as  to  obtain  a  steady  stream  of  bubbles 

through  B.    Col- 

lect  two   jars   of 

the  gas    by   dis- 

placement of  air, 

then  transfer  the 

delivery  -  tube   to 

a  bottle  contain- 

ing   water,    in 

which     the     gas 

will  dissolve  and 

form  a   solution. 

When  the  evolu- 

tion   of     gas    is 

nearly     finished, 

disconnect  the  apparatus  and  add  50  c.c.  of  water  to  the  flask  A  ' 

raise  the  contents  to  boiling,  and  filter  the  liquid  into  a  porcelain 

dish.      Concentrate  the  blue  solution  to  about  half  its  bulk  by 

evaporation,  and  set  it  aside  to  crystallize.     Blue  crystals  of  copper 

sulphate,  CuS04,  5H20,  are  obtained. 

The  principal  change  is  expressed  by  the  equation  : 

Cu  4-  2H2S04  =  CuS04  +  2H20  +  S02. 
Properties  of  Sulphur  Dioxide.—  The  gas  collected  in 

the  jars  is  colourless  and  transparent,  and  has  the  odour  of 
burning  sulphur.  When  a  burning  candle  is  lowered  into 
one  of  the  jars,  the  flame  disappears,  and  the  gas  itself  does 
not  burn.  A  solution  of  blue  litmus  is  immediately  turned 
red  when  poured  into  the  jar.  When  the  other  jar  is  inverted 
over  water  and  the  plate  removed,  the  water  rises  rapidly 
up  the  jar  to  take  the  place  of  the  dissolved  gas.  The 
solution,  made  as  described  above,  smells  strongly  of  the  gas, 
and  is  considered  to  contain  sulphurous  acid,  H2S03,  formed 
by  the  reaction  between  the  gas  and  water,  thus  : 


The  pure  acid  has  not  been  isolated  ;    it  is  only  known  in 
solution,  but  it  forms  salts  called  sulphites. 


96  METALLURGICAL  CHEMISTRY 


METALS  AND  NITRIC  ACID. 

Many  important  changes  depend  upon  the  readiness  with 
which  nitric  acid  parts  with  oxygen  to  bodies  capable  of  com- 
bining with  that  element,  and  an  interesting  series  of  reactions 
may  be  brought  about.  It  may  be  remarked  at  the  outset 
that  metals  are  more  readily  soluble  in  nitric  acid  than  in 
either  of  the  two  acids  already  dealt  with.  The  metals  may 
be  roughly  divided  into  three  classes  with  respect  to  the  action 
of  nitric  acid  upon  them. 

1 .  The   metal   is   converted   directly   into   its   oxide :    Sn 
and  Sb. 

2.  The  metal  is  converted  into  a  soluble  nitrate :  Cu,  Mg, 
Hg,  Pb,  Fe,  Ag,  Zn. 

3.  The  metal  is  not  attacked  by  the  acid  :  Au  and  Pt. 

The  solid  non-metals  also  are  more  or  less  readily  converted 
into  oxides,  which  then  combine  with  water  to  form  the 
corresponding  acids. 

Tin — EXP.  56. — Heat  a  clean  porcelain  crucible  over  the  Bunsen 
flame.  When  cold,  weigh  it  carefully,  and  weigh  in  it  0*5  gram  of 
tin.  Add  5  c.c.  of  dilute  nitric  acid  (1  to  1)  from  a  pipette.  Put  a 
watch-glass  over  the  crucible,  and  heat  it  carefully  on  the  sand-bath 
or  iron  plate.  When  the  action  is  finished,  remove  the  watch-glass, 
and  if  any  white  particles  adhere  to  it,  rinse  them  into  the  crucible 
with  a  fine  jet  of  water  from  the  wash-bottle.  Now  place  the 
crucible  on  a  pipe-clay  tripod  (see  Exp.  53),  and  evaporate  the 
contents  slowly  to  dryness.  Then  put  the  crucible  in  a  pipe-clay 
triangle  on  an  iron  tripod  stand,  and  heat  it  carefully  with  the 
Bunsen  flame.  Hold  the  burner  in  the  hand,  move  the  flame 
about,  and  be  ready  to  remove  it  if  the  residue  shows  signs  of 
spurting.  When  risk  of  loss  is  over,  heat  the  crucible  strongly 
for  ten  minutes,  and  then  for  a  minute  or  two  over  the  foot  blow- 
pipe flame.  If  a  muffle  is  available,  it  can  be  used  for  the  final 
heating.  When  cold,  weigh  the  crucible  and  oxide. 

EXAMPLE. — Weight  of  crucible  +  oxide  =  11*770  grams 


„      =  0-635  gram 
Then  0'635  -0*5  =  0135  =  weight  of  oxygen  absorbed. 


METALS  AND  ACIDS  97 

0*5 
To  find  the  formula  of  the  compound  we  have  :  vy^  =  0*0042 

and  -T*—  =  0*0084.  Therefore  the  ratio  of  the  number  of  atoms 
ID 

of  tin  to  the  number  of  atoms  of  oxygen  in  the  molecule  is 

0-0042 

0-0084  =  *  : 

Antimony— EXP.  57. — Use  0-5  gram  of  powdered  antimony,  and 
5  c.c.  of  dilute  acid  (1  in  3;.  Follow  the  instructions  given  for  the 
last  experiment.  The  final  heating  must  be  done  with  the  crucible 
open,  for  if  a  lid  is  used,  reducing  gases  may  get  into  the  crucible 
and  cause  loss  by  volatilization. 

EXAMPLE. — Weight  of  metal  used  =  0'5  gram;  weight  of  oxygen 
absorbed  =  0-138  gram.  Use  the  data  obtained  to  find  the  fomiula 
of  the  oxide.  Sb  =  120. 

When  metals  dissolve  in  nitric  acid  their  nitrates  are 
formed.  They  are,  however,  readily  decomposed  on  being 
heated,  and,  in  most  cases,  are  converted  into  the  oxides  of 
the  contained  metals.  There  are  two  notable  exceptions  to 
this :  When  nitrate  of  silver  is  heated  it  is  completely 
decomposed,  and  a  residue  of  metallic  silver  is  obtained ; 
when  nitrate  of  mercury  is  heated  it  is  first  reduced  to  the 
oxide,  which  on  further  heating  splits  up  into  mercury  and 
oxygen. 

During  the  dissolution  of  a  metal  by  nitric  acid  in  an  open 
vessel  a  red  gas  appears  to  be  given  off  as  long  as  any  of  the 
metal  remains.  If,  then,  the  solution  is  evaporated  to  dryness, 
and  the  solid  residue  strongly  heated,  a  further  evolution  of 
red  gas  is  noticed.  The  final  residue  is  found  to  be  the  oxide 
of  the  metal,  and  its  weight  compared  with  that  of  the  metal 
used  indicates  the  weight  of  oxygen  absorbed. 

Iron — EXP.  58. — Put  0'5  gram  of  clean  iron  wire  into  a  weighed 
porcelain  crucible,  and  add  5  c.c.  of  dilute  nitric  acid  (1  to  1). 
Dissolve,  evaporate,  and  heat  the  residue  as  described  in  Exp.  56. 
Allow  the  crucible  to  cool,  and  reweigh  it.  If  the  wire  is  thick, 
it  should  be  first  flattened  under  the  hammer  and  rubbed  with 
emery-cloth  before  being  weighed  and  dissolved. 

7 


98  METALLUKGICAL  CHEMISTKY 

EXAMPLE.  —  Weight  of  metal  used  =  0'5  gram  ;  weight  'of  oxygen 
absorbed  =  0-215  gram. 


Then       =  0-0089,  and  =  0-0134.     Therefore 

or,  taking  the  nearest  whole  numbers,  the  ratio  is  2  :  3,  and  the 
formula  of  the  oxide  is  Fe2O3. 

Mercury  —  EXP.  59.  —  Use  about  1  gram  of  mercury  and  5  c.c. 
of  dilute  acid  (1  to  2).  Proceed  as  in  Exp.  58.  After  the  evapora- 
tion, to  obtain  the  white  residue  of  nitrate,  the  crucible  must  be 
heated  very  carefully  by  holding  the  Bunsen  burner  in  the  hand 
and  moving  it  about.  The  white  colour  of  the  residue  gradually 
changes  to  red,  and  the  action  is  finished  when  its  colour  is 
uniformly  red.  It  must  not  become  dark-coloured  during  the 
heating,  or  loss  of  metal  and  oxygen  will  take  place.  Eeweigh 
the  crucible  when  cold. 

EXAMPLE.  —  Weight  of  mercury—  1'07  grams;  increase—  0'075 
gram. 

Find  the  formula  of  the  oxide  from  the  data  thus  obtained. 
Hg  =  200. 

Silver  —  EXP.  60.  —  Use,  0*5  gram  of  silver  and  5  c.c.  of  dilute 
acid  (1  to  1).  A  residue  =  0*5  gram  of  silver  is  obtained. 

Similar  experiments  may  be  made  with  zinc,  magnesium, 
lead,  and  bismuth.  The  result  with  the  last-named  metal 
leads  to  the  formula  BLOo. 


THE  GASES  EVOLVED  WHEN  METALS  AND  NITRIC  ACID 
REACT. 

It  will  be  useful  to  consider,  as  far  as  can  be  done  in  an 
elementary  way,  the  nature  of  the  gases  given  off  when  metals 
dissolve  in  nitric  acid.  It  has  already  been  noticed  that  a  red 
gas  appears  to  be  given  off  almost  invariably  when  the  dis- 
solution takes  place  in  an  open  vessel.  But  when  the  experi- 
ment is  made  in  a  closed  vessel,  it  is  found  that  the  gas 
actually  escaping  from  the  liquid  is  not  of  constant  composi- 
tion for  all  metals.  In  fact,  for  the  same  metal  even  the 
volume  and  composition  may  vary  with — (1)  the  concentration 
of  the  acid ;  (2)  the  temperature  at  which  the  reaction  is  pro- 


METALS  AND  ACIDS  99 

ceeding;  (3)  the  depth  of  the  layer  of  acid  liquid  above  the 
dissolving  metal,  and  through  which  the  escaping  gas  has  to 
pass.  It  also  varies  with  the  nature  of  the  metal. 

During  the  changes  which  nitric  acid  undergoes  several 
oxides  of  nitrogen,  and  even  nitrogen  itself,  may  be  liberated. 
The  hydrogen  of  the  acid  is  almost  invariably  converted  into 
water  by  union  with  oxygen.  Under  exceptional  circum- 
stances it  may  be  converted  into  ammonia,  NH3,  by  union 
with  nitrogen,  or  even  liberated  in  the  free  state. 

There  are  three  well-known  and  easily-obtainable  oxides  of 
nitrogen :  nitrous  oxide,  N20  ;  nitric  oxide,  NO  ;  and  nitric 
peroxide,  N02,  or  N204. 

Nitrous  Oxide,  N20,  is  a  colourless  gas  with  a  faint  odour 
and  sweet  taste.  It  does  not  change  when  mixed  with  air, 
and  is  somewhat  soluble  in  water.  If  a  glowing  splint  is  put 
into  a  jar  of  this  gas,  the  splint  bursts  into  flame ;  but  when 
feebly-burning  sulphur  is  brought  into  contact  with  it,  the 
flame  disappears.  This  test  distinguishes  nitrous  oxide  from 
oxygen.  It  is  used  by  dentists  to  produce  insensibility  to 
pain  for  a  short  time.  Its  popular  name  is  "  laughing  gas." 

EXP.  61. — Fit  up  apparatus  similar  to  that  shown  in  Fig.  23,  but 
use  an  8-ounce  bottle,  so  as  not  to  have  a  large  volume  of  air  for  the 
gas  to  displace,  and  arrange  a  small  wash- bottle,  nearly  full  of 
strong  nitric  acid,  in  the  path  of  the  gas  to  absorb  any  nitric  oxide 
which  may  pass  over.  Put  an  excess  of  zinc  into  the  bottle,  and  add 
a  small  quantity  of  water  through  the  thistle  funnel  to  just  cover 
the  bottom.  Place  the  bottle  in  a  dish  of  water  to  prevent  a  too 
rapid  rise  in  temperature,  and  add  dilute  nitric  acid  (1  to  1),  a  little 
at  a  time,  to  keep  up  a  steady  stream  of  bubbles.  Collect  two  jars 
of  the  gas,  over  warm  water  preferably.  Remove  the  plate  from 
one  of  the  jars;  no  change  in  colour  is  observed.  Plunge  a  glowing 
splint  into  the  jar ;  the  splint  bursts  into  flame.  Heat  a  little  sulphur 
in  a  deflagrating  spoon  until  it  shows  a  small  flame,  and  plunge  it 
into  a  jar  of  the  gas  ;  the  flame  disappears.  Heat  the  sulphur  until 
it  is  burning  vigorously,  and  put  it  into  the  jar  again ;  it  continues 
to  burn,  and  with  a  brighter  flame  than  when  burning  in  the  air. 

Nitrous  oxide  supports  combustion  more  vigorously  than 
air  does,  but  the  burning  body  must  have  a  sufficient  tempera. 

7—2 


100  METALLUEGICAL  CHEMISTRY 

ture  to  decompose  the  gas  and  abstract  its  oxygen,  if  the 
combustion  is  to  continue. 

Nitric  Oxide,  NO,  is  a  colourless  gas,  which  is  converted 
very  rapidly  into  nitric  peroxide  when  mixed  with  air  or 
oxygen.  The  gas  thus  formed  has  a  red  colour,  so  that  the 
change  is  easily  observed,  and  it  is  commonly  used  to  dis- 
tinguish oxygen  from  nitrous  oxide,  which  does  not  react  with 
nitric  oxide.  Thus,  the  addition  of  nitric  oxide  is  a  test  for 
oxygen,  and  the  addition  of  oxygen  is  a  test  for  nitric  oxide 
in  mixtures  of  gases.  Nitric  oxide  is  only  slightly  soluble  in 
water,  and  does  not  support  the  combustion  of  ordinary  burning 
bodies. 

EXP.  62.  —  Eepeat  the  last  experiment,  using  thin  sheet  copper  in 
the  place  of  zinc.  The  bottle  need  not  be  cooled,  and  the  wash- 
bottle  must  be  omitted.  Collect  two  jars  of  the  gas,  and  remove 
them  from  the  trough.  Collect  one  tall  cylinder  about  two-thirds 
full  of  the  gas,  and  leave  it  standing  over  water.  Eemove  the  plate 
from  one  of  the  jars  ;  a  red  gas  forms  in  the  mouth  of  the  jar. 
Plunge  a  burning  candle  into  the  jar  :  the  flame  disappears.  Just 
ignite  a  small  piece  of  phosphorus  in  a  deflagrating  spoon  and 
plunge  it  rapidly  into  a  jar  of  the  gas  :  the  flame  disappears.  Allow 
the  phosphorus  to  burn  vigorously,  and  return  it  to  the  jar  :  it  burns 
more  vigorously  in  the  nitric  oxide  than  in  air.  Decant  half  the 
volume  of  oxygen  into  the  jar  standing  over  water  :  grip  the  side  of 
the  jar  in  the  hand  while  transferring  the  oxygen  to  it.  A  red  gas 
forms,  and  dissolves  rapidly,  the  water  rising  in  the  jar  after  each 
addition  of  oxygen,  and  the  jar  becomes  quite  warm  in  the  hand. 

Nitric  oxide,  although  not  a  supporter  of  combustion  in  the 
ordinary  sense,  will  give  up  oxygen  to  bodies  which  have  a 
temperature  sufficiently  high  to  effect  its  decomposition. 

The  reaction  between  nitric  oxide  and  oxygen  is  expressed 
by  the  equation  : 


Heat  is  developed  by  the  combination  of  the  two  gases. 

Nitric  Peroxide,  N02,  is  a  red  gas,  very  soluble  in  water, 
with  which  it  reacts,  and  forms  an  acid  solution  containing 
nitrous  and  nitric  acids.  The  peroxide  itself  is  not  an 


METALS  AND  ACIDS  101 

acid-forming  oxide.  The  change  is  expressed  by  the  equa- 
tion : 

2N02  +  H20  =  HN03  +  HN02. 

Nitric         Nitrous 
acid.  acid. 

The  formula  N02  does  not  represent  the  composition  of  nitric 
peroxide  molecules  under  all  conditions.  At  a  low  tempera- 
ture the  composition  is  more  nearly  represented  by  N204.  If 
the  temperature  is  allowed  to  rise,  the  colour  of  the  gas 
deepens,  and  the  complex  molecules  of  N204  are  gradually 
decomposed  into  the  simpler  molecules  of  N02.  At  a 
moderately  high  temperature  the  composition  of  the  gas  is 
represented  by  N02.  At  a  still  higher  temperature  the 
whole  of  the  gas  is  decomposed  into  nitric  oxide  and  oxygen, 
and  the  colour  entirely  disappears.  As  the  temperature  falls 
the  inverse  changes  take  place.  This  is  a  good  example  of 
"  dissociation." 

Nitric  peroxide  readily  gives  up  part  of  its  oxygen  to 
bodies  eager  to  combine  with  that  element.  This  is  taken 
advantage  of  in  the  manufacture  of  sulphuric  acid,  during 
which  nitric  oxide  is  made  to  act  as  a  carrier  of  oxygen  from 
the  air  to  sulphur  dioxide  for  its  conversion  into  the  trioxide, 
thus  : 

NO  +  0  =  N02,  and  then  N02  +  S02  =  S03  +  NO. 

These  reactions  are  repeated  as  long  as  nitric  oxide,  sulphur- 
dioxide,  and  oxygen  are  present  in  the  acid  chambers. 

When  silver  dissolves  into  nitric  acid,  both  nitric  peroxide 
and  nitric  oxide  are  liberated. 

The  principal  reactions  of  zinc,  copper,  and  silver  with  nitric 
acid  may  be  represented  by  equations  : 

1.  4Zn  +  10HN03  =  4Zn(N03)2  +  5H20  +  N20. 

2.  3Cu  +  8HN03  =  3Cu(N03)2  +  4H2O  +  2NO. 
3. 


The  nitrate  of  the  metal  is  formed  in  each  case. 

It  is  usually  stated  that  hydrogen  gas  cannot  be  obtained 


102 


METALLURGICAL  CHEMISTEY 


by  the  action  of-  a  metal  on  nitric  acid ;  but  if  magnesium  is 
substituted  for  zinc  in  Exp.  61,  the  bottle  kept  cool,  and  the 
acid  dilute,  the  greater  part  of  the  evolved  gas  is  found  to  be 
hydrogen.  But  if  the  acid  is  too  strong,  and  the  temperature 
is  allowed  to  rise,  the  proportion  of  hydrogen  in  the  evolved 
gas  is  considerably  diminished. 

EFFECTS  OF  DILUTION  AND  TEMPERATURE. 

The  effects  of  a  variation  in  the  strength  of  nitric  acid,  and 
in  the  temperature  at  which  the  reaction  takes  place  on  the 
volume  of  the  escaping  gas,  are  interesting,  and  instructive 

experiments  may  be 
made  with  the  appa- 
ratus shown  in  Fig.  25. 
A  is  a  test-tube  fitted 
with  a  bung  and  angle- 
tube.  B  is  a  small  tap 
funnel.  C  is  a  fixed 
delivery-tube  con- 
nected with  the  angle- 
tube  from  A  by  a 
piece  of  sound  rubber 
tube,  and  passing 
under  the  mouth  of  the  graduated  cylinder  D.  If  it  is 
desired  to  keep  the  test-tube  A  cool  during  an  experiment, 
it  is  immersed  in  a  beaker  containing  cold  water ;  or,  when 
the  beaker  is  away,  the  test-tube  can  be  heated  by  a  Bunseii 
flame. 

The  conduct  of  an  experiment  is  simple.  When  the  parts 
of  the  apparatus  are  in  position,  water  is  poured  into  B  and 
the  tap  opened.  The  water  running  from  B  into  A  displaces 
the  air  in  A  and  the  delivery-tube  connected  with  it.  The 
displaced  air  collects  in  D,  and  its  volume  can  be  measured. 
Thus  the  volume  of  air  in  the  apparatus  at  the  beginning  of 
an  experiment  is  known,  if  the  volume  of  the  added  acid  is 


FIG.  25. 


METALS  AND  ACIDS  103 

subtracted  from  the  total  volume  obtained  as  above.  The 
water  is  then  emptied  out,  and  the  necessary  quantity  of  acid 
added  to  A  ;  the  coil  of  weighed  metal  is  dropped  in,  and  the 
bung  quickly  replaced.  If  the  tube  is  to  be  kept  cool,  the 
beaker  of  water  is  placed  in  position.  When  the  metal  is 
completely  dissolved,  water  is  poured  into  B,  the  tap  opened, 
and  the  whole  of  the  gas  driven  from  A  to  D.  The  volume 
of  gas  in  D  is  then  measured  by  depressing  the  cylinder  in 
the  trough  until  the  level  of  the  water  is  the  same  inside  and 
outside,  and  then  taking  the  reading  on  the  graduated  scale. 
Half  the  volume  of  oxygen  is  then  passed  into  D  by  decant- 
ing the  gas  from  a  graduated  cylinder.  When  the  red  gas 
formed  by  the  reaction  between  the  nitric  oxide  and  oxygen 
is  all  dissolved,  and  the  residue  is  colourless,  the  cylinder  is 
depressed,  and  the  volume  of  the  residual  gas  measured. 
Sufficient  data  are  thus  obtained  for  roughly  finding  the  com- 
position of  the  evolved  gas. 

When  copper  is  the  dissolving  metal,  nitric  oxide,  NO,  is 
the  principal  constituent  of  the  liberated  gas.  Now,  it  has 
been  shown  already  that  when  this  gas  is  mixed  with  oxygen 
the  red  gas,  nitrogen  peroxide,  N02,  is  'formed  and  it  is  easily 
proved  that  nitric  oxide  combines  with  exactly  half  its  own 
volume  of  oxygen  during  the  change.  The  reaction  is  shown 
by  the  equation  : 


4  vols.     2  vols. 

Iii  the  experiments  made  to  obtain  the  following  results,  strips 
of  thin  sheet  copper  of  uniform  size,  weighing  exactly  1  gram, 
were  used,  together  with  15  c.c.  of  the  acid.  In  this  way  the 
volume  of  air  in  the  apparatus  was  kept  the  same  in  each 
experiment.  Now,  it  must  be  borne  in  mind  that  one-fifth  of 
the  enclosed  air  is  oxygen,  which  will  combine  with  twice  its 
own  volume  of  nitric  oxide. 

EXAMPLE.  —  Volume  of  air  in  the  apparatus  after  adding  the 
acid  =  50  c.c.  Total  volume  of  gas  in  D  after  dissolution  of  the 
metal,  and  displacement  of  gas  in  the  apparatus  by  water  run  in 


104 


METALLURGICAL  CHEMISTRY 


from  B  =  158  c.c. ;  volume  of  oxygen  added  =  80  c.c. ;  volume  of 
residual  gas  in  D  =  76  c.c. 

Then  158 +  80 -76 -162  =  the  volume  of  NO  and  02, 
which  has  disappeared,  and,  since  two-thirds  of  this  is  NO 

162  x  2 

—  =108  =  the  volume  of  nitric  oxide  in  the  collected  gas. 

o 

But  20  c.c.  of  NO  are  also  accounted  for  by  the  10  c.c.  of 
oxygen  in  the  enclosed  air  at  the  beginning  of  the  experiment. 
Therefore,  the  total  volume  of  NO  liberated  =  128  c.c. 

Also  the  residual  gas  contains  40  c.c.  of  nitrogen  from 
the  enclosed  air,  and  26  c.c.  of  added  oxygen.  Therefore, 
76-40-26  =  10  c.c.  of  gas  other  than  nitric  oxide  liberated 
from  the  acid. 

Therefore,   the    NO    in    the    liberated    gas  =  — ^-^ —  = 

92-7  per  cent. 

The  following  table  shows  the  results  of  two  experiments 
conducted  as  described  above.  It  demonstrates  clearly  that 
the  volume  and  nature  of  the  gas  liberated  during  the  reaction 
between  nitric  acid  and  a  particular  metal  depends  upon  the 
conditions  under  which  the  reaction  takes  place.  One  gram  of 
copper  was  used  for  each  experiment. 

It  may  be  remarked  that  nitric  oxide  is  soluble  in  the  nitric 
acid  solution,  and  that  its  solubility  decreases  as  the  tempera- 
ture increases.  This  would  account  largely  for  the  difference 
in  the  volumes  of  the  gases  liberated  during  the  two  experi- 
ments. The  changes  taking  place  between  nitric  acid  and 
metals  deserve  more  attention  than  they  have  yet  received ; 
but  the  subject  is  a  difficult  one,  and  the  experiments  described 
above  only  give  a  rough  indication  of  the  method. 


STRENGTH  OF 
ACID. 

VOLUME 

OF 

VOLUME 

OF 

OXYGEN 

VOLUME 

OF 

RESIDUAL 

PERCENT- 
AGE OF 
NITRIC 

TEMPERATURE. 

ACID. 

WATER. 

GAS. 

ADDED. 

GAS. 

OXIDE. 

1 

1 

158  c.c. 

80  C.C. 

76  c.c. 

927 

Allowed  to  rise 

1 

1 

106    „ 

52    „ 

84    „ 

77-5 

Kept  constant 

METALS  AND  ACIDS  105 

ACTION  OF  ACIDS  UPON  COMMERCIAL  IRON. 

The  different  varieties  of  commercial  iron — known  as  cast 
iron,  steel,  and  wrought  iron — owe  their  variations  in  properties 
very  largely  to  the  presence  of  carbon  in  the  metal.  This 
element  may  be  present,  not  only  in  varying  proportions,  but 
also  in  different  conditions,  and  thus  cause  a  wide  variation  in 
the  properties  of  the  metal. 

For  our  present  purpose  carbon  may  be  said  to  exist  in 
iron  in  two  distinct  forms.  As  Graphite  it  is  mechanically 
mixed  in  the  form  of  crystalline  scales  with  the  particles  of 
the  iron.  These  scales  are  readily  seen  on  looking  at  the 
fractured  surface  of  a  piece  of  grey  pig-iron  with  a  pocket 
lens.  As  Carbide  it  is  combined  with  part  of  the  iron  in  the 
form  of  iron  carbide,  Fe3C,  which  cannot  be  observed  without 
special  preparation  of  the  surface  of  the  metal,  and  its  examina- 
tion under  the  microscope. 

Cast  iron  contains  on  the  average  3 '5  per  cent,  of  carbon, 
which  exists  largely  as  graphite  in  grey  iroir,  largely  as  carbide 
in  white  iron,  or  in  both  forms  in  mottled  iron.  The  amount 
of  carbon  in  steel  varies  considerably,  and  may  be  as  high  as 
1*8  per  cent.  When  the  steel  is  in  its  normal  or  annealed 
state,  the  carbon  is  in  the  form  of  carbide.  Wrought  iron 
may  contain  up  to  0  2  per  cent,  of  carbon,  which  is  present  as 
carbide.  On  account  of  these  variations  the  common  acids, 
though  able  to  dissolve  the  metal  readily,  give  different  results 
with  different  varieties. 

Hydrochloric  and  dilute  sulphuric  acids  act  similarly,  but 
the  action  varies  with  the  condition  of  the  carbon.  The  rate 
of  dissolution  increases  with  the  temperature.  Graphite 
separates  as  such,  and  being  insoluble  in  the  acid  liquid,  settles 
out,  and  forms  a  black  deposit  on  the  bottom  of  the  vessel  in 
which  the  dissolution  is  taking  place.  The  carbide  is  decom- 
posed by  the  acid,  assisted  by  the  newly  liberated  hydrogen, 
and  the  freed  carbon  combining  with  hydrogen  forms  hydro- 
carbons, which  mix  with  the  escaping  gas,  and  give  to  it  a 


106  METALLUEGICAL  CHEMISTKY 

very  disagreeable  odour.  The  presence  of  carbide  carbon  is 
detected  in  this  way. 

Dilute  nitric  acid  acts  differently  as  regards  the  carbide 
carbon,  which  in  this  case  passes  into  the  solution  in  combina- 
tion with  oxygen  and  hydrogen.  The  compound  formed 
imparts  a  distinct  colour  to  the  solution.  Under  similar  con- 
ditions the  depth  of  the  colour  varies  with  the  proportion  of 
carbon  present.  Graphite,  however,  is  set  free,  and  separates 
out  in  the  same  way  as  with  the  other  acids. 

The  other  elements  present  in  commercial  iron  are  silicon, 
phosphorus,  sulphur,  and  manganese,  and  the  purity  of  the 
metal  depends  upon  the  proportions  present,  which  may  vary 
from  traces  to  several  per  cents.  Each  one  has  its  own  parti- 
cular influence  upon  the  properties  of  the  metal,  and  has  to  be 
reckoned  with  by  the  practical  man.  During  the  dissolution 
of  the  metal  in  acids  they  undergo  various  changes.  The 
silicon  is  oxidized  to  silica  and  left  as  a  solid  residue,  inter- 
mixed with  any  graphite  that  may  be  present.  Phosphorus 
and  sulphur  are  oxidized  to  phosphoric  and  sulphuric  acids  by 
moderately  strong  nitric  acid,  and  pass  into  solution.  With 
hydrochloric  and  sulphuric  acids  the  sulphur  escapes  entirely 
as  sulphuretted  hydrogen  gas,  and  the  phosphorus  partially  as 
phosphoretted  hydrogen.  The  iron  and  any  manganese  which 
may  be  present  dissolve  in  and  form  salts  of  the  acids  used. 

ACTION  OF  ACIDS  ON  METALLIC  ALLOYS. 

Alloys  are  bodies  containing  two  or  more  metals,  which  may 
be  present  under  different  conditions.  Sometimes  the  relation 
between  them  is  so  intimate  that  the  alloy  may  be  regarded 
as  a  chemical  compound.  Sometimes  the  body  is  simply  a 
solidified  solution  of  its  constituent  metals. 

As  a  rule,  neither  hydrochloric  nor  dilute  sulphuric  acid 
will  dissolve  these  bodies,  and  the  more  intimate  the  relation 
between  their  constituents,  the  more  difficult  they  are  to  dis- 
solve. Dilute  nitric  acid  is  the  best  general  solvent,  and  is 
commonly  used.  If  both  metals  in  an  alloy  dissolve  indi- 


METALS  AND  ACIDS  107 

vidually  in  the  acid,  then  the  alloy  is  also  soluble ;  but  if  one 
of  the  metals  is  insoluble,  it  is  left  as  a  residue  of  oxide  or 
metal,  as  the  case  may  be.  A  copper-zinc  alloy  (brass)  is 
completely  dissolved.  A  copper-tin  alloy  (bronze)  is  partially 
dissolved,  and  a  residue  of  white  tin  oxide  is  obtained.  A 
gold-silver  alloy  containing  more  than  2J  parts  of  silver  to 
1  part  of  gold  is  partially  dissolved,  and  a  residue  of  metallic 
gold  left. 

The  last  two  reactions  furnish  the  means  of  separating  tin 
from  copper,  and  gold  from  silver. 

EXP.  63. — Put  about  1  gram  each  of  brass  and  bronze  into 
separate  test-tubes  ;  add  10  c.c.  of  dilute  nitric  acid  to  each,  and 
heat  the  tube,  if  necessary,  until  all  action  ceases.  Repeat  the 
experiment,  using  dilute  sulphuric  and  hydrochloric  acids  in  place 
of  the  nitric  acid. 

When  brass  dissolves,  copper  and  zinc  nitrates  are  formed  ; 
and  the  partial  dissolution  of  bronze  results  in  the  formation 
of  copper  nitrate  and  hydrated  tin  oxide. 

ELECTRO-CHEMICAL  ACTION. 

An  important  result  of  the  dissolution  of  a  metal  in  an  acid 
has  now  to  be  considered.  It  has  already  been  noticed  that 
when  the  action  is  brisk,  the  vessel  in  which  it  is  taking  place 
gets  quite  hot.  It  may,  then,  be  said  that  heat  is  one  of  the 
products  of  the  chemical  change ;  and  under  ordinary  circum- 
stances, the  amount  of  heat  evolved  is  just  as  definite  as  the 
volume  of  gas  liberated,  or  the  weight  of  salt  formed.  This 
is  easy  to  prove  with  suitable  apparatus,  but  it  is  outside 
ordinary  elementary  work,  and  must,  therefore,  be  here  taken 
for  granted. 

Pure  zinc  is  scarcely  attacked  by  cold  dilute  sulphuric  acid, 
but  if  a  piece  of  copper  is  placed  in  contact  with  the  zinc  in 
the  dilute  acid,  the  action  begins,  and  gas  is  liberated  from  the 
surface  of  the  copper.  It  would  thus  appear  at  first  sight 
as  though  the  copper  were  dissolving,  but  it  soon  becomes 
evident  that  this  is  not  so.  The  solution  remains  colourless, 


108  METALLUEGICAL  CHEMISTKY 

which  would  not  be  the  case  if  copper  sulphate  were  being 
formed.  Also,  the  zinc  gradually  disappears,  but  the  copper, 
if  weighed  before  being  put  in,  and  weighed  again  after  being 
taken  out  of  the  solution,  is  found  to  be  unchanged.  Clearly, 
then,  the  contact  of  the  copper  with  the  zinc  causes  the  latter 
to  dissolve  more  readily,  but  does  not  otherwise  affect  the 
chemical  change.  And  the  solution  gets  hot,  too,  so  that  heat 
is  evolved  as  usual. 

If  ordinary  sheet  zinc  is  attacked  by  the  dilute  acid,  it 
should  be  cleaned  by  dipping  it  in  moderately  strong  acid, 
and  then  coated  with  mercury.  Zinc  so  treated  is  said  to  be 
amalgamated,  and  acts  with  the  dilute  acid  in  the  same  way 
as  the  pure  metal.  Silver,  platinum,  or  gas-carbon  may  be 
used  in  the  place  of  copper,  to  set  up  the  action. 

EXP.  64. — Put  a  piece  of  pure  zinc  or  the  amalgamated  metal 
into  a  test-tube.  Add  a  few  cubic  centimetres  of  dilute  sulphuric 
acid  (1  to  20),  and  allow  the  tube  to  stand  for  a  minute  or  two. 
Note  that  no  action  is  going  on.  Drop  a  piece  of  sheet  copper  on 
to  the  zinc.  The  action  commences  at  once,  and  soon  bubbles  of 
gas  are  freely  liberated  from  the  copper.  The  solution  does  not 
change  colour,  and  the  zinc  gradually  dissolves. 

A  silver  coin,  a  piece  of  platinum  foil,  or  a  carbon  rod  may 
be  substituted  for  the  copper  with  similar  results. 

If  long  strips  of  copper  and  zinc  are  only  partly  immersed 
in  the  acid  liquid,  and  the  ends  outside  are  made  to  touch,  the 
same  result  is  obtained.  Zinc  dissolves,  and  hydrogen  gas  is 
liberated  from  the  copper,  even  though  the  immersed  portions 
are  some  distance  apart.  Also  the  ends  of  the  plates  may  be 
joined  by  a  metal  wire  of  considerable  length  with  similar 
results.  But  directly  the  connection  is  severed  the  action 
ceases. 

EXP.  65.  — Cut  strips  of  thick  sheet  zinc  and  sheet  copper,  about 
6  inches  long  by  1  inch  wide  ;  immerse  the  zinc  strip  about  three- 
quarters  the  way  up  in  moderately  strong  sulphuric  acid  for  a  few 
seconds,  and  then  rub  a  little  mercury  all  over  the  cleansed  surface 
on  both  sides.  Fix  the  plates,  one  on  each  side  of  a  strip  of  wood 
1  inch  wide,  so  that  their  ends  project  a  little  above  the  upper 
surface  of  the  wood.  This  is  easily  done  by  driving  short  tacks 


METALS  AND  ACIDS 


109 


through  the  plates  into  the  wood,  but  not  in  the  same  line,  so  that 
there  will  be  no  danger  of  their  meeting  in  the  wood.  Put  the 
plates  into  a  beaker  con- 
taining dilute  sulphuric 
acid,  which  should 
nearly  cover  the  amal- 
gamated part  of  the 
zinc  plate.  Connect  the 
plates  by  forcing  the 
clean  ends  of  a  copper 
wire  between  the  plates 
and  the  wood.  Observe 
that  the  action  com- 
mences as  soon  as  the 
connection  is  made. 
Make  a  similar  pair, 
but  with  two  copper  A  FIG.  26.  B 

plates  instead  of  copper 

and  zinc.  Put  this  pair  into  a  solution  of  copper  sulphate,  and 
connect  the  ends  by  copper  wires,  as  shown  in  Fig.  26.  Allow  the 
action  to  go  on  for  a  few  minutes  ;  then  take  out  the  copper  plates, 
and  examine  them. 


The  copper  sulphate  solution  makes  the  connection,  and  the 
action  goes  on  in  A  in  much  the  same  way  as  when  the  con- 
nection is  made  direct  by  a  single  wire.  The  copper  plate  in 
B,  which  is  connected  directly  with  the  zinc  plate  in  A,  is 
found  to  be  coated  with  freshly  deposited  copper,  and  the 
longer  the  action  is  allowed  to  proceed,  the  thicker  the  deposit 
becomes.  In  fact,  if  the  two  plates  are  separately  weighed 
before  and  after  immersion  in  B,  one  is  found  to  increase  and 
the  other  to  decrease  in  weight  to  exactly  the  same  extent. 
If  the  connections  are  reversed,  copper  is  deposited  on  the 
other  plate.  Evidently  some  kind  of  influence  is  transferred 
from  A  to  B,  which  carries  the  copper  from  one  plate  to  the 
other,  and  its  direction  depends  upon  the  way  in  which  the 
connection  is  made.  It  seems  to  move  from  Cu  in  A  to  Cul 
in  B,  across  the  solution  to  Cu2,  and  back  to  Zn  in  A.  This 
is  called  an  electric  current. 

If  two  or  three  cells  similar  to  A  have  their  alternate  copper 
and  zinc  plates  connected  by  short  thick  copper  wires,  and 


110  METALLUEGICAL  CHEMISTKY 

the  terminal  copper  and  zinc  are  connected  with  carbon  rods 
or  platinum  plates  in  B,  copper  is  still  deposited  on  the  plate 
connected  with  the  terminal  Zn.  But  in  this  case  nothing 
dissolves  from  the  other  plate,  and  the  copper  sulphate  is 
gradually  decomposed,  oxygen  being  liberated  from  the  plate 
connected  with  the  terminal  Cu.  Now,  if  the  amount  of  heat 
developed  by  the  dissolution  of  zinc  in  the  cell  during  the  de- 
composition of  the  copper  sulphate  is  determined,  it  is  found 
to  fall  considerably  short  of  the  quantity  which  would  be 
developed  by  the  dissolution  of  the  same  weight  of  zinc  in  the 
ordinary  way.  The  missing  heat  appears  in  the  form  of 
electrical  energy,  which  is  carried  round  the  system  by  the 
current,  and  determines  the  decomposition  of  the  sulphate. 

If  a  very  long  thin  iron  wire  is  used  to  connect  the  two 
plates  in  A  (Fig.  26),  nearly  the  whole  of  the  energy  developed 
by  the  dissolution  of  the  zinc  is  transferred  to  the  wire.  It 
appears  first  as  electrical  energy,  is  then  carried  through  the 
wire  by  the  current,  and  finally  transformed  into  heat  by  the 
resistance  of  the  wire  to  the  passage  of  the  current.  With  a 
short  thick  copper  wire  for  connections,  nearly  the  whole  of 
the  heat  appears  in  A. 

The  principle  demonstrated  in  the  last  experiment  is  the 
one  upon  which  the  electro-deposition  of  metals  depends,  and 
it  is  put  to  a  variety  of  uses.  A  large  number  of  metals  can 
be  so  deposited,  and  it  is  only  necessary  that  the  surface  upon 
which  the  deposit  takes  place  should  be  a  conductor  of 
electricity.  The  common  plan  is  to  deposit  a  metal  which  has 
specific  properties,  such  as  gold  and  silver,  on  account  of  their 
lustre,  and  nickel  in  virtue  of  its  hardness.  It  is  even  possible 
to  deposit  two  metals  together,  so  as  to  form  an  alloy.  Thus 
the  deposition  of  copper  and  zinc  to  form  a  deposit  of  brass  is 
common.  The  purification  of  metals  by  the  use  of  electrical 
energy  is  carried  out  on  the  large  scale.  Electrolytic  copper 
is  a  very  pure  metal,  which  is  refined  by  this  process,  and  is 
largely  used  for  electrical  purposes,  and  for  making  high-class 
copper  alloys. 


METALS  AND  ACIDS  HI 

Electrolysis  is  the  term  used  to  denote  the  process  through 
which  a  metallic  salt  or  oxide  passes  when  it  is  decomposed 
by  an  electric  current.  The  body  undergoing  the  change  is 
usually  dissolved  in  water;  but  some  solids  which  can  be 
melted  without  undergoing  rapid  decomposition  are  electro- 
lized  in  the  molten  condition.  The  decomposing  compound 
is  called  the  electrolyte,  and  the  facts  of  electrolysis  point 
strongly  to  a  decomposition  of  at  least  some  of  the  molecules 
of  the  compound  by  its  dissolution  in  the  solvent  liquid  alone. 
These  molecules,  however,  are  only  dissociated,  for  if  the 
solvent  is  removed  by  evaporation  the  original  compound  is 
recovered  intact.  According  to  this,  when  copper  sulphate  is 
the  electrolyte  it  seems  that  an  appreciable  number  of  its 
molecules  are  dissociated  into  Cu  and  S04.  The  Cu  re- 
presents an  elementary  atom  or  simple  radicle,  and  the  S04  a 
group  of  atoms  or  a  compound  radicle,  which  appears  to  have 
properties  similar  to  those  of  an  elementary  atom.  It  must 
be  clearly  understood  that  these  radicles  are  not  supposed  to 
exist  separately  in  the  same  way  as  an  ordinary  element  or 
compound.  The  existence  of  the  one  predicates  the  co-exist- 
ence of  the  other  in  the  solution  of  the  salt.  They  are  known 
as  ions,  and  the  function  of  the  current  is  supposed  to  be  to 
direct  them  towards  the  electrodes  by  which  it  enters  and 
leaves  the  solution.  The  first  is  the  anode,  to  which  the 
anions  move,  and  the  second  the  cathode,  to  which  the  cathions 
move.  In  the  electrolysis  of  CuS04,  the  S04  groups  are  the 
anions  and  the  Cu  atoms  the  cathions.  When  copper  elec- 
trodes are  used  the  Cu  ions  simply  deposit  on  the  cathode, 
and  lose  their  ionic  condition,  while  the  S04  ions  deposit  on 
the  anode  and  lose  their  ionic  condition  by  abstracting  copper 
from  the  plate  itself.  But  when  platinum  electrodes  are  used 
the  S04  ions  attack  water  molecules  in  the  proximity  of  the 
anode,  form  H2S04  molecules,  and  set  oxygen  free,  which 
collects  on  the  anode  plate.  The  change  with  the  copper 
anode  is  in  effect  a  transport  of  copper  from  one  electrode  to 
the  other ;  with  the  platinum  anode  it  is  equivalent  to  the 


112  METALLURGICAL  CHEMISTRY 

decomposition  of  copper  oxide  into  its  elements.  Much  more 
electrical  energy  is  absorbed  in  the  second  case  than  in  the 
first;  it  is  used  up  in  effecting  the  chemical  change  in  the 
electrolyte. 

This  is  shown  by  the  equations  : 


The  decomposition  of  water  by  the  current  (p.  27)  when 
regarded  from  this  point  of  view,  is  really  effected  by  the 
electrolysis  of  a  solution  of  sulphuric  acid  : 


This  is  supported  by  the  fact  that  pure  water  will  not  convey 
the  current. 

The  chief  electrolytes  are  solutions  of  mineral  acids  and 
metallic  salts,  molten  salts  and  oxides,  or  solutions  of  oxides 
in  molten  salts. 

Voltaic  Battery.—  The  simple  cell  described  in  Exp.  65 
falls  off  in  strength  very  rapidly,  owing  to  the  accumulation  of 
a  film  of  hydrogen  on  the  copper  plate.  This  offers  consider- 
able resistance  to  the  passage  of  the  current  through  the  cell, 
and  also  sets  up  a  back  pressure  which  diminishes  the  main 
current.  This  is  avoided  in  some  constant  cells  by  oxidizing 
the  hydrogen  to  water,  and  thus  preventing  its  deposition  on 
the  plate.  The  Bunsen  cell  is  made  up  as  follows  :  A  circular 
earthenware  vessel  containing  dilute  sulphuric  acid,  in  which 
is  immersed  a  sheet  of  amalgamated  zinc  bent  round  into  the 
form  of  a  hollow  cylinder  ;  a  porous  earthenware  pot  contain- 
ing strong  nitric  acid,  in  which  is  immersed  a  carbon  rod. 
The  porous  pot  stands  inside  the  zinc  cylinder,  and  binding 
screws  are  attached  to  the  zinc  and  carbon  plates.  The  liquids 
soak  into  the  wall  of  the  porous  pot  from  both  sides,  and  a 
conducting  medium  extends  between  the  two  plates.  When 
the  binding  screws  are  joined  by  a  conductor,  a  much  stronger 
and  more  constant  current  than  can  be  obtained  by  a  simple 
cell  passes  through  the  wire.  The  function  of  the  nitric  acid  is 


METALS  AND  ACIDS  113 

to  prevent  the  deposition  of  hydrogen  on  the  carbon  plate,  by 
oxidizing  it  to  water.  By  linking  several  of  these  cells  to- 
gether in  series,  a  current  sufficiently  powerful  for  ordinary 
experiments  is  obtained. 

SUMMARY. 

When  a  metal  dissolves  in  hydrochloric  or  dilute  sulphuric 
acid,  the  reaction  is  simple,  for  hydrogen  only  is  liberated,  and 
a  salt  of  the  metal  and  acid  formed.  The  volume  of  hydrogen 
set  free  depends  upon  the  weight  of  the  metal  used,  and  is 
different  for  the  same  weight  of  different  metals.  With  hot 
strong  sulphuric  acid  the  reaction  is  more  complex,  and 
sulphur  dioxide  is  the  principal  gas  liberated. 

Nitric  acid  dissolves  most  of  the  common  metals,  and  gives 
some  very  complicated  reactions.  The  products  of  its  decom- 
position are  varied,  several  oxides  of  nitrogen,  hydrogen, 
ammonia,  and  nitrogen  itself  being  liberated  under  different 
conditions.  The  nitrates  of  the  metals  are  formed  in  most 
cases,  but  with  tin  and  antimony  the  oxides  of  the  metals 
are  produced.  The  nitrates  are  readily  decomposed  by  heat, 
and,  except  in  one  or  two  cases,  a  residue  of  the  oxides  is 
obtained. 

The  experiments  on  the  oxidation  of  metals  in  air  (Chap.  II.) 
and  by  nitric  acid  give  interesting  comparative  results.  If  a 
fixed  weight  of  oxygen  is  taken,  and  the  weights  of  different 
metals  which  would  combine  with  this  fixed  weight  are  cal- 
culated from  the  results  of  experiments,  an  interesting  table 
can  be  constructed,  as  follows  : 

Combining  weight  of  oxygen  =  8. 


METAL. 

PARTS  BY  WEIGHT. 

METAL. 

PARTS  BY  WEIGHT. 

Magnesium     .  .  . 
Iron     ... 
Tin       
Copper 

12  0 
18-7 
30-0 
31-5 

j  Zinc     ... 
Bismuth 
Mercury 
|  Lead    ... 

...  !               32-5 
...  i               69-0 
lOO'O 
103-5 

METALS  AND  ACIDS  115 

The  numbers  are  readily  obtained,  for  in  Exp.  69  : 

0*251  gram  of  oxygen  combines  with  1  gram  of  copper 

/.     1          ,          „          would  combine  with 


.-.    8    grams      „  „  „ 

or  copper 

QUESTIONS. 

1.  Describe  an  experiment  by  which  the  volume  of  hydrogen 
liberated  by  the  dissolution  of  1  gram  of  iron  in  hydrochloric 
acid  can  be  determined. 

2.  What   changes   take   place    (a)  when  zinc  dissolves  in 
dilute  sulphuric  acid  ;  (b)  when  copper  and  the  strong  acid  are 
heated  together  ? 

3.  How  would  you  find  the  weight  of  magnesium  sulphate 
formed  when  O5  gram  of  the  metal  dissolves  in  dilute  sulphuric 
acid? 

4.  Give  a  short  description  of  the  general  reactions  between 
metals  and  nitric  acid. 

5.  Name  the  three  common  oxides  of  nitrogen,  and  briefly 
describe  them. 

6.  What  are  the  general  effects  of  dilution  and  temperature 
on  the  changes  taking  place  between  nitric  acid  and  a  given 
metal  1 

7.  Which   common   acid  is   the   best  general   solvent  for 
metals  and  alloys,  and  why  ? 

8.  Describe  a  simple  voltaic  cell,  and  say  how  you  would 
obtain  an  electric  current  with  it. 

9.  WThat  do  you  understand  by  the  term  "electrolysis"? 
Give  an  example  to  illustrate  your  answer. 


8—2 


CHAPTER  VIII 

CHEMICAL  EQUIVALENTS  AND  ATOMIC  WEIGHTS  OF 
COMMON  METALS  AND  NON-METALS 

THE  student  is  now  familiar  with  the  fact  that  elements 
combine  together  to  form  compounds,  and  is  also  aware  of 
the  very  definite  character  of  these  compounds,  for  he  has 
brought  about  various  combinations,  and  has  seen  the  extra- 
ordinary changes  effected  in  the  properties  of  the  elements 
between  which  these  combinations  have  taken  place.  Also,  in 
carrying  out  the  experiments  on  the  dissolution  of  metals  in 
acids,  he  has  been  directed  to  the  conclusion  that  there  is  a 
definite  relation  between  the  weight  of  metal  dissolved  and 
the  volume  of  gas  liberated  during  its  dissolution.  It  is  also 
fully  recognised  that  the  weight  of  a  given  body  of  gas  is  just 
as  definite  as  its  volume.  This  is  assumed  in  Chap.  VIL, 
together  with  the  known  weight  of  a  litre  of  hydrogen,  which 
is  used  to  calculate  the  weights  of  the  volumes  of  gas  evolved 
in  the  experiments  described ;  but  it  is  possible  to  use  a  more 
direct  method,  if  the  pure  gas  only  is  allowed  to  escape  from 
the  apparatus,  and  the  loss  in  weight  due  to  this  escape  is 
noted. 

The  object  of  the  following  experiments  is  to  make  it  clear 
that  certain  proportional  numbers  or  weights  are  inseparably 
associated  with  the  elements  when  they  are  entering  into  com- 
bination with  each  other,  or  are  being  expelled  from  their 
compounds ;  and  that  these  numbers,  which  are  the  results  of 
actual  experiments,  are  connected  in  a  simple  way  with  the 
atomic  weights  of  the  elements  to  which  they  refer.  The 

116 


EQUIVALENTS  AND  ATOMIC  WEIGHTS 


117 


results  we  obtain  may  be  only  approximate ;  but  if  they  are 
sufficiently  accurate  to  give  a  clear  idea  of  the  facts  they  are 
intended  to  illustrate,  their  object  will  be  attained.  Experi- 
ments carried  out  by  experienced  workers  using  accurate 
apparatus  fully  justify  the  deductions  drawn  from  them,  which 
may  therefore  be  accepted  with  confidence. 

To  FIND  THE  WEIGHT  OF  HYDROGEN  EXPELLED  FROM  AN 

ACID   BY    A   KNOWN   WEIGHT   OF  MAGNESIUM. 

EXP.  66. — Fit  up  the  apparatus  shown  in  Fig.  27.     A  is  a  100  c.c. 
flask  fitted  with  a  two-hole  bung.     A  straight  tube,  B,  fitted  with 
a  cap  to  make  it  air  tight,  passes  through  one  hole,  and  a  bent  tube, 
C,  through  the  other.     The  other  end  of  C  passes  through  a  bung 
in  the  U-tube,  D.    The  bend  of  the  U-tube 
is  filled  with   glass-wool,  and  each  limb 
with  small  pieces  of  fused  calcium  chloride, 
and  glass-wool  on  the  top.     The  free  limb 
of  the  U-tube  is  fitted  with  a  bung  and 
capped  tube.     The  actual  apparatus  should 
weigh  about  70  grams.    The  caps  are  easily 
made  by  pushing  short  pieces  of  glass  rod 
into  pieces  of  rubber  tube  about  twice  their 
length. 

Add  10  c.c.  of  dilute  H2S04  (1  in  12)  to 
the  flask  A,  wipe  the  apparatus,  and  let  it 
stand  for  a  time  near  the  balance.  Weigh 
accurately  about  0'2  gram  of  Mg  ribbon, 
and  make  it  into  a  coil,  which  will  pass 
readily  down  the  neck  of  A,  and  stand  just 
clear  of  the  bung  when  dropped  in.  Put  the  apparatus  and  weighed 
magnesium  on  the  scale-pan,  and  weigh  them  together.  Remove 
the  cap  from  the  tube  D.  Take  the  bung  out  of  A,  drop  in  the 
metal  coil,  and  quickly  replace  the  bung.  As  the  metal  dissolves 
the  coil  gradually  descends  until  complete  dissolution  is  effected. 
By  this  arrangement  the  evolution  of  gas  is  rendered  uniform  and 
not  too  rapid. 

Place  the  apparatus  near  the  balance  until  it  is  nearly  cold,  and 
then  slowly  aspirate  about  a  litre  of  dry  air  through  it.  The  air 
is  dried  by  passing  through  a  U-shaped  tube  fitted  in  the  same 
manner  as  D,  and  attached  to  B  by  a  rubber  tube.  Eeplace  the 
caps,  and  reweigh  the  apparatus.  The  loss  in  weight  is  due  to  the 
escape  of  dry  hydrogen. 

The  aspiration  of  air  through  the  apparatus  is  of  consider- 


FIG.  27. 


118 


METALLURGICAL  CHEMISTRY 


able  importance,  as  a  large  portion  of  the  liberated  hydrogen 
remains  in  the  apparatus  when  the  dissolution  is  finished. 
Any  form  of  aspirator  may  be  used.  In  Fig.  28  the  inlet  of 
the  flask  is  connected  with  the  drying-tube  B,  and  the  outlet 
of  the  drying-tube,  through  which  the  gas  escapes,  with  the 


FIG.  28. 

aspirator  A.  The  tap  C  is  opened  carefully,  and  about  1 
litre  of  water  allowed  to  slowly  escape.  The  hydrogen  jis  thus 
replaced  by  air,  and  the  atmosphere  in  the  apparatus  is  now 
similar  to  what  it  was  at  the  beginning  of  the  experiment. 

EXAMPLE. — Weight  of  magnesium  taken  =  0*201  gram. 

,,       „   apparatus  +  magnesium  =79-296  grams. 
„       „  „         -hydrogen      =  79«279      „ 

,,       ,,  hydrogen  liberated  =  0'017  gram. 

Simplify  this  result  by  calculating  the  weight  of  magnesium 
required  to  liberate  1  part  by  weight  of  hydrogen  : 

0-017  gram  of  hydrogen  is  replaced  by  0*201  of  magnesium. 


o^rr 

Or,  11 '82  parts  of  magnesium  are  equivalent  to   1  part  of 
hydrogen. 


EQUIVALENTS  AND  ATOMIC  WEIGHTS  119 

To  FIND  THE  WEIGHT  OF  HYDROGEN  EXPELLED  FROM  AN 

ACID   BY  A   KNOWN   WEIGHT   OF   ZlNC. 

EXP.  67.— Repeat  the  last  experiment  with  about  0'75  gram  of 
thin  zinc-foil  and  10  c.c.  of  dilute  sulphuric  acid  (1  in  6).  The 
stronger  acid  is  recommended,  as  the  action  is  not  so  violent  as  with 
the  magnesium. 

EXAMPLE.— Weight  of  zinc  taken  =       0'743  gram. 

,,       ,,  apparatus  +  zinc  =     77*688  grams. 

„       „          „         -hydrogen=     77'665       „ 

„       ,,  hydrogen  liberated      =       0'023  gram. 
Then,  as  before  : 

O023  gram  of  hydrogen  is  replaced  by  0*743  of  zinc. 

0-743     00  o    r    • 
'  =  32-3  of  zinc. 


Or,  32 '3  parts  of  zinc  are  equivalent  to  1  part  of  hydrogen. 

These  results  may  be  used  as  starting-points  in  the  deter- 
mination of  the  equivalent  weights  of  other  elements. 

To  FIND  THE  WEIGHT  OF  OXYGEN  WHICH  WILL  COMBINE 

WITH   A   KNOWN   WEIGHT    OF   MAGNESIUM. 

EXP.  68. — Heat  a  large  porcelain  crucible  and  lid  (No.  2)  over  the 
Bunsen  flame,  and  set  it  aside  to  cool.  Take  a  piece  of  magnesium 
ribbon  18  inches  long  and  weighing  about  0'25  gram ;  coil  it  round 
a  piece  of  thin  glass  rod,  and  cut  the  long  coil  into  six  short  ones. 
Weigh  the  crucible  and  lid,  put  in  the  coils,  and  weigh  again.  Put 
the  crucible  over  a  good  Bunsen  flame  on  a  pipeclay  triangle,  and 
slope  the  lid  a  little  so  as  to  give  a  narrow  air  space  on  one 
side.  Heat  the  crucible  strongly  for  fifteen  minutes  without  raising 
the  lid.  Then  raise  the  lid  carefully,  and  notice  if  there  are  any 
signs  of  burning.  If  not,  remove  the  flame,  and  break  up  the 
residue  with  a  clean  glass  rod,  being  careful  to  brush  back  into  the 
crucible  any  particles  which  may  adhere  to  the  rod.  Continue  the 
heating  without  the  lid  for  fifteen  minutes  longer.  Allow  the 
crucible  to  cool,  and  weigh  it.  To  be  quite  certain  that  the  action  is 
finished,  the  crucible  should  be  heated  again  for  a  few  minutes, 
cooled^  and  reweighed  to  see  if  the  weight  is  constant. 


120  METALLURGICAL  CHEMISTKY 

EXAMPLE.— Weight  of  crucible  +  magnesium  =  33-789  grams. 

-33-544     „ 


=  0-245  gram. 
„    .  ,,  after  heating  =  33-951  grams. 


„  before     „          33'789 


Increase     =  0'162  gram. 

The  increase  is  due  to  oxygen  absorbed  from  the  air. 
Take  12  the  nearest  whole  number  to  the  equivalent  value 
found  for  magnesium,  and  find  the  combining  number  for 
oxygen  : 

0*245gramof  magnesium  combines  with  0"  162  of  oxygen. 

0-162 


^245 


•'•    12        „  „  ,  ^^—  =  7-93  grams. 

Then  8  may  be  taken  for  the  combining  proportion  of 
oxygen. 

To  FIND  THE  WEIGHT  OF  COPPER  WHICH  WILL  COMBINE 

WITH   A   GIVEN   WEIGHT   OF   OXYGEN. 

EXP.  69.  —  Heat  a  No.  1  porcelain  crucible  over  the  Bunsen  flame. 
Allow  it  to  cool  ;  weigh  it,  and  weigh  in  it  1  gram  of  finely-divided 
copper  reduced  from  the  oxide.  Put  the  crucible  in  a  fire-clay 
dish,  and  place  the  dish  in  a  moderately  hot  munie.  Continue  the 
heating  for  forty-five  minutes  ;  remove  the  crucible  ;  cool,  and 
weigh  it.  Repeat  the  heating  for  ten  minutes,  and  reweigh  to  see  if 
the  weight  is  constant. 

Grams. 

EXAMPLE.—  Weight  of  crucible  +  copper  after  heating   12-851 

before    ,  12-600 


Increase      =     0-251 
Weight  of  copper  taken  =  1  gram. 

Now,  0*251  gram  of  oxygen  combines  with  1  gram  of  copper ; 

Q 

.-.  8  grams  of  oxygen  combine  with  -          =31 -88  of  copper. 

0'251 

31-5  may  be  taken  as  the  combining  proportion  of  copper. 


EQUIVALENTS  AND  ATOMIC  WEIGHTS  121 

NOTE. — The  reduced  copper  for  the  experiment  is  readily 
prepared  by  reducing  powdered  copper  scale  by  coal  gas  or 
hydrogen.  The  common  commercial  oxide  is  too  impure  to 
give  a  good  result. 


To  FIND  THE  WEIGHT  OF  COPPER  EXPELLED  FROM  ITS 
COMPOUNDS  BY  A  KNOWN  WEIGHT  OF  ZINC. 

EXP.  70. — Dissolve  about  3  grams  of  crystallized  copper  sulphate, 
CuS04,  5H20,  in  50  c.c.  of  water  ;  filter  the  solution,  if  necessary, 
into  a  small  beaker.  Weigh  accurately  about  0*5  gram  of  thin  zinc- 
foil,  and  transfer  it  to  the  sulphate  solution.  Heat  the  solution 
gently,  and  stir  occasionally  with  a  glass  rod  until  the  zinc  is  com- 
pletely dissolved.  This  is  determined  by  breaking  up  the  precipitated 
copper  with  the  end  of  the  glass  rod.  While  the  dissolution  is 
proceeding  fold  two  filter-papers  and 
put  them  in  a  filter-funnel  in  the  air- 
oven,  or  over  the  sand-bath,  to  dry. 
When  dry  place  one  in  each  pan  of  the 
balance,  and  counterpoise  them,  cutting 
the  apex  off  the  heavier  one  and  then 
snipping  the  one  or  the  other  at  the  top 
until  they  counterpoise.  Then  put  the 
complete  filter  into  the  one  which  has 
had  the  apex  cut  off,  and  open  it  out  so 
as  to  make  a  single  filter  with  four  folds 
of  paper  on  each  side.  Put  it  into  the  Fio.  29. 

funnel,  and  when  ready  filter  the  sul- 
phate solution.  Transfer  the  whole  of  the  precipitated  copper  to 
the  filter,  and  wash  it  thoroughly  until  every  trace  of  the  excess  of 
copper  sulphate  has  been  washed  through.  Use  a  wash-bottle  in 
which  the  water  can  be  heated,  as  the  washing  is  effected  more 
rapidly  with  hot  than  with  cold  water.  Place  the  filter  funnel  and 
filter  in  a  cone  over  the  sand-bath  or  in  the  air-oven.  Be  careful 
not  to  overheat  the  filter,  or  the  finely- divided  copper  will  absorb 
oxygen,  and  too  high  a  result  will  be  obtained.  When  the  copper  is 
thoroughly  dry,  separate  the  two  filters  and  put  them  on  the  scale- 
pans  in  the  same  order  as  when  they  were  counterpoised ;  add 
weights  until  they  counterpoise  again.  The  difference  gives  the 
weight  of  the  copper.  Eepeat  the  drying  for  a  few  minutes  to  see  if 
the  weight  is  constant. 

By  using  two  filters  as  described  loss  of  copper  is  avoided,  and 
by  putting  them  both  in  the  funnel  any  loss  of  weight  due  to  wash- 
ing is  the  same  in  both. 


122  METALLURGICAL  CHEMISTRY 

EXAMPLE.— Weight  of  zinc  taken  =     0-528  gram. 

,,      ,,  copper  obtained  =     0'509     „ 

Then  using  the  equivalent  for  zinc  found  by  Exp.  67,  calculate 
the  equivalent  for  copper. 

0"528  gram  of  zinc  displaced  0-509  gram  of  copper  ; 

0-509 
"    ^528 

0-509  x  32-3 
.\32-3          „  „  „    -——_  —=31-47  of  copper. 

U  DZO 

Compare  the  results  of  the  last  two  experiments. 


To  FIND  THE  WEIGHT  OF  SILVER  EXPELLED  FROM  ITS 
COMPOUNDS  BY  A  KNOWN  WEIGHT  OF  COPPER. 

EXP.  71.  —  Dissolve  about  a  gram  of  thin  sheet  silver  with  5  c.c. 
of  strong  sulphuric  acid  in  a  small  beaker.  Put  the  beaker  on  a 
sand-bath  in  the  fume-chamber,  and  heat  it  carefully  until  the 
metal  is  dissolved.  Allow  it  to  get  cold,  and  then  pour  down  the 
side  about  40  c.c.  of  distilled  water.  A  white  crystalline  precipitate 
of  silver  sulphate  forms,  which  dissolves  again  on  boiling  the 
solution.  Weigh  accurately  about  0-25  gram  of  thin  copper-foil,  put 
it  into  the  hot  silver  solution,  and  allow  to  stand  until  the  copper 
is  completely  dissolved.  The  dissolution  may  be  hastened  by 
keeping  the  solution  hot,  but  it  is  better  to  let  it  stand  in  the  cold. 
Counterpoise  two  filter-papers  as  described  in  Exp.  70.  Filter 
off  the  precipitated  silver,  thoroughly  wash  and  dry  it.  Weigh, 
dry  again,  and  reweigh  until  the  weights  are  constant.  The  filtrate 
contains  silver,  which  can  be  recovered  by  immersing  a  strip  of 
copper  in  it. 

EXAMPLE.  —  Weight  of  copper  taken      =  0'251 
,,       ,,  silver  obtained    =0'860 

Then,  taking  31  "5  as  the  equivalent  for  copper  : 

0-251  gram  of  copper  displaced  0-860  gram  of  silver  ; 

0-860 


.-.  31-5  grams  „         „         „  =  10>71  of  silver. 

108  may  be  taken  as  the  equivalent  of  silver. 


EQUIVALENTS  AND  ATOMIC  WEIGHTS  123 

TO   FIND   THE   WEIGHT   OF   CHLORINE  WHICH   WILL    COMBINE 
WITH  A   KNOWN   WEIGHT   OF   SILVER. 

EXP.  72. — Weigh  accurately  about  0*5  gram  of  fine  silver; 
transfer  it  to  a  small  beaker,  and  add  5  c.c.  of  dilute  nitric  acid 
(1  to  1)  ;  heat  the  beaker  gently  on  the  hot  plate  until  the  metal  is 
dissolved  ;  make  up  the  solution  to  about  100  c.c.  with  distilled 
water ;  raise  it  to  boiling,  and  add  a  little  dilute  hydrochloric  acid  ; 
stir  well  with  a  glass  rod  until  the  white  curdy  precipitate  collects 
together  ;  add  a  few  drops  more  of  the  acid  to  see  if  the  pre- 
cipation  is  complete.  Add  the  acid  carefully,  so  as  to  have  only 
a  slight  excess  in  the  solution  when  the  whole  of  the  silver  is 
precipitated.  Dry  and  counterpoise  two  filter-papers  as  in  Exp.  70. 
Filter  off  the  silver  chloride,  and  wash  it  thoroughly  with  hot 
water.  Dry  the  chloride  carefully,  separate  the  filters,  and  weigh. 
Kepeat  the  drying  and  weighing  until  a  constant  weight  is 
obtained. 

EXAMPLE. — Weight  of  silver  chloride  obtained  =  0'67  5  gram. 

taken=0-509      „ 
,,         „  chlorine  found  =  0-166      ,, 

0*509  gram  of  silver  combines  with  0-166  gram  of  chlorine ; 

0-166 
"          "          "          "     0-509      " 

0-166x108      f  85-22  grams 
)8  grams  „          „          „          „       Q.509       -|  of  chiorine. 

35-5  may  be  taken  as  the  equivalent  of  chlorine. 

If  potassium  bromide  and  potassium  iodide  solutions  are 
used  to  precipitate  known  weights  of  silver,  in  the  same  way 
as  described  in  the  last  experiment,  the  combining  proportions 
of  bromine  and  iodine  can  be  determined  as  readily  as  that  of 
chlorine. 

If  known  weights  of  sodium  chloride  and  potassium  bromide 
are  precipitated  with  excess  of  silver  nitrate,  and  the  precipi- 
tates collected  and  weighed,  the  combining  proportions  of 
sodium  and  potassium,  compared  with  chlorine  and  bromine 
respectively,  are  obtained. 


124  METALLUKGICAL  CHEMISTRY 

SUMMARY  OF  RESULTS. 

The  experiments  by  which  the  results  given  in  the  previous 
examples  were  obtained  were  made  with  special  care,  so  as  to 
furnish  reliable  standards  for  the  student  to  work  to.  There- 
fore, the  beginner  must  not  expect  too  much  from  his  own 
initial  experiments,  and  should  be  satisfied  if  his  results  are 
sufficiently  near  the  true  ones  to  give  him  confidence  in  the 
general  principles  which  are  under  investigation,  and  in  the 
results  which  are  stated  as  the  outcome  of  many  most  accurate 
investigations. 

It  should  be  borne  in  mind  that  if  two  or  more  distinct 
methods  are  used  to  determine  the  chemical  equivalent  of  a 
particular  element,  the  results  should  be  practically  the  same. 
This  is  illustrated  in  the  determination  of  the  equivalent  of 
copper :  (a)  by  its  combination  with  oxygen  ;  (b)  by  its  pre- 
cipitation by  zinc. 

The  common  starting-point  is  hydrogen,  because  it  has  the 
lowest  equivalent  of  any  known  element.  It  therefore  fur- 
nishes a  very  simple  standard ;  but  any  other  element  might 
be  used.  Berzelius,  who  was  the  first  to  furnish  a  well- 
defined  list  of  the  equivalents  or  combining  proportions  of 
elements,  adopted  0  =  100  for  his  standard.  But  it  is  readily 
seen  that  such  a  standard  would  be  somewhat  cumbersome  in 
its  application. 

The  chemical  equivalents  or  combining  proportions  of 
elements  are  experimental  numbers,  and  do  not  depend  upon 
any  theory  of  the  constitution  of  compounds,  other  than  that 
they  contain  elements  in  chemical  combination.  No  difficulty 
would,  however,  be  experienced  in  assigning  symbols  to  these 
combining  proportions,  and  in  representing  compounds  with 
them ;  but  formulae  of  compounds  derived  in  this  way  would 
not  agree  in  many  cases  with  those  found  by  the  application 
of  the  theory  of  atoms.  It  is  therefore  necessary  to  consider 
the  question  from  another  point  of  view,  and  endeavour  to 
trace  a  relation  between  chemical  equivalents  and  atomic 
weights. 


EQUIVALENTS  AND  ATOMIC  WEIGHTS 


125 


EELATION  OF  CHEMICAL  EQUIVALENTS  TO  ATOMIC  WEIGHTS. 


ELEMENT. 

SYMBOL. 

i 

EQUIVALENT. 

ii 

ATOMIC  WEIGHT. 

RATIO  OF  i  :  ii 

Hydrogen 

H 

1-0 

i-o 

1 

Chlorine 

Cl 

35-22 

35-5 

1 

Silver 

Ag 

10771 

1077 

1 

Oxygen 

0 

7-93 

16-0 

2 

Magnesium 

Mg 

11-82 

24-0 

2 

Copper 

Cu 

31-47 

63-0 

2 

Zinc  ... 

Zn 

32-3 

65-0 

2 

Bismuth 

Bi 

69-21 

207-5 

3 

Tin*  ... 

Sn 

29-4 

117-4 

4 

Atomic  Weights.  —  Certain  numbers,  termed  atomic 
weights,  have  been  used  in  connection  with  elements  in  pre- 
ceding chapters,  and  a  reference  to  the  table  given  above 
indicates  that  the  combining  proportion  of  an  element  as  deter- 
mined by  experiment  does  not  always  agree  with  its  atomic 
weight  as  commonly  used.  Although  the  relation  between  the 
two  is  a  simple  one,  it  requires  some  explanation  if  the  beginner 
is  to  use  atomic  weights  with  the  confidence  they  deserve.  It 
may  be  stated  that  the  determination  of  the  combining  pro- 
portion of  an  element  is  always  the  first  step  in  the  fixing  of 
its  atomic  weight,  and  that  the  other  considerations  are  also 
based  on  experimental  grounds.  What  these  considerations 
are  depends  upon  the  nature  of  the  element  and  its  compounds ; 
so  that  a  method  which  is  very  useful  in  dealing  with  some 
elements  may  entirely  fail  with,  others.  Often  two  or  more 
methods  may  be  applied  to  the  same  element,  and  if  their 
results  agree,  the  deductions  drawn  from  them  are  all  the 
more  convincing.  The  method  of  deducing  an  atomic  weight 
may  depend  upon  either  the  properties  of  the  element  itself 
or  on  those  of  its  compounds.  The  important  methods  used 
for  fixing  the  atomic  weights  of  elements  when  their  combining 
proportions  are  known  will  now  be  briefly,  but  sufficiently, 
described,  to  enable  the  student  to  get  a  general  grasp  of  the 

.       *  See  Exp.  56. 


126  METALLUEGICAL  CHEMISTRY 

subject ;  but  for  fuller  details  larger  works  on  chemistry  must 
be  consulted. 

Vapour  Density. — A  number  of  elements  and  compounds 
are  either  gaseous  at  the  temperature  of  the  atmosphere,  or 
are  converted  into  gas  when  raised  to  temperatures  within  the 
range  of  accurate  experimental  work.  The  weights  of 
measured  volumes  of  these  gases  or  vapours  can  be  determined, 
and  compared  with  the  weights  of  equal  volumes  of  hydrogen. 
This  gives  the  vapour  density  of  the  element  or  compound  in 
the  gaseous  state. 

There  are  several  methods  employed  for  the  determination 
of  vapour  density ;  but  a  brief  outline  of  a  comparatively 
simple  one  will  be  sufficient  for  our  present  purpose.  Suppose 
the  compound  under  consideration  to  be  a  liquid  which  is 
easily  converted  into  a  combustible  vapour.  A  glass  flask 
with  its  neck  drawn  off  to  a  narrow  tube  is  weighed,  and  an 
excess  of  the  liquid  introduced.  The  flask  is  then  immersed 
in  a  bath,  the  temperature  of  which  is  well  above  the  boiling- 
point  of  the  liquid,  when  a  considerable  volume  of  the  vapour 
escapes  and  carries  with  it  the  air  which  originally  filled  the 
flask.  When  the  air  is  all  expelled,  and  the  liquid  is  com- 
pletely converted  into  vapour,  the  flask  is  full  of  it  at  the 
temperature  of  the  bath  and  the  pressure  of  the  outside  atmo- 
sphere. If  the  jet  of  vapour  is  allowed  to  burn  as  it  issues 
from  the  neck,  the  end  of  the  vaporization  is  indicated  by  a 
sudden  drop  in  the  flame.  The  end  of  the  neck  is  now  sealed 
up  by  softening  the  glass  in  the  blow-pipe  flame,  the  flask 
removed  from  the  bath,  and  weighed  when  cold.  The  end  of 
the  neck  is  then  broken  under  water,  the  flask  entirely  filled, 
and  weighed  again  full  of  water.  The  weight  of  the  water  in 
grams  gives  the  volume  of  the  flask,  for  a  gram  of  water 
measures  a  cubic  centimetre.  The  weights  obtained,  together 
with  the  corrections  to  reduce  the  volume  of  the  vapour  at 
the  temperature  and  pressure  of  the  experiment  to  its  volume 
at  N.T.P.  (see  Appendix),  and  for  the  expansion  of  the  flask, 


EQUIVALENTS  AND  ATOMIC  WEIGHTS  127 

give  the  weight  and  volume  of  the  vapour  under  examination. 
It  is  then  easy  to  compare  this  weight  with  the  weight  of  an 
equal  volume  of  hydrogen  under  the  same  conditions  of 
temperature  and  pressure.  It  may,  then,  be  stated  that  the 
vapour  density  of  an  element  or  a  compound  which  can  be 
converted  into  vapour  without  decomposition  or  dissociation 
is  an  experimental  number  ;  and  full  advantage  may  be  taken 
of  the  fact  in  helping  to  fix  atomic  weights. 

If  it  is  assumed  that  the  bulk  of  1  part  by  weight  of 
hydrogen  represents  the  volume  of  an  atom  of  the  element,  it 
may  be  chosen  as  the  unit  of  comparison  for  other  atomic  and 
molecular  volumes.  Twice  this  bulk  of  hydrogen  represents 
a  molecule  of  the  gas,  and,  by  Avogadro's  law  (see  p.  43), 
any  other  molecule  in  the  gaseous  state  fills  the  same  volume. 
From  the  definition  of  the  term  atom,  a  molecule  of  any  kind 
cannot  contain  less  than  an  atom  of  an  element. 

So  that  the  smallest  weight  of  an  element  ever 
found  in  a  molecular  volume  of  the  vapour  of  any  of 
its  compounds  may  be  taken  as  the  atomic  weight  of 
that  element. 

Further,  it  appears  that  vapour  densities  furnish  a  direct 
method  of  determining  molecular  weights ;  for  since  it  is 
assumed  that  a  molecule  of  hydrogen  contains  two  atoms,  its 
weight  is  twice  that  of  a  single  atom,  so  that  the  weight  of  a 
molecule  of  any  other  gas  must  be  proportional  to  twice  its 
vapour  density.  If,  then,  the  atomic  weights  of  the  elements 
in  a  compound  are  known,  and  its  vapour  density  can  be 
determined,  a  definite  molecular  weight  is  readily  assigned  to 
it.  For  example,  the  smallest  weight  of  oxygen  ever  found 
in  two  volumes  of  any  of  its  volatile  compounds  is  16;  the 
vapour  density  of  water  is  9;  and,  therefore,  its  molecular 
weight  is  18 ;  so  that  a  molecule  of  water  contains  an  atomic 
weight  of  oxygen,  and  2  atomic  weights  of  hydrogen  and  its 
formula  is  therefore  H90. 

The  majority  of  the  metals  cannot  be  vapourized  at  workable 


128  METALLURGICAL  CHEMISTRY 

temperatures,  nor  do  they  furnish  volatile  compounds,  so  that 
the  application  of  the  above  principle  to  the  fixing  of  their 
atomic  weights  is  limited.  The  problem  must  in  their  case  be 
solved  by  other  means,  as  described  below. 

Isomorphism. — Metals  can  be  arranged  more  or  less 
definitely  into  groups,  such  that  the  members  of  any  one 
group  can  replace  each  other  in  solid  compounds  without  alter- 
ing the  crystalline  form  of  the  compounds.  These  metals  and 
their  crystallized  compounds  are  said  to  be  isomorphous,  and 
their  atoms  are  regarded  as  fulfilling  similar  functions  in  the 
molecules  of  the  compounds.  If,  then,  the  atomic  weight  of 
one  metal  in  a  group  has  been  fixed  by  any  other  method,  the 
atomic  weights  of  the  other  members  of  the  same  group  can 
be  deduced.  Thus  the  chemical  equivalent  of  aluminium  is  9, 
and  it  forms  only  one  oxide,  in  which  the  proportion  Al  :  0  is 
18  :  16 ;  if  it  is  inferred  that  the  formula  of  this  oxide  is  A10, 
the  atomic  weight  of  the  metal  must  be  taken  Al  =  18.  But 
from  other  considerations  the  formula  of  the  red  oxide  of  iron 
is  given  as  Fe2O3,  and  the  oxide  of  aluminium  is  isomorphous 
with  it,  so  that  the  more  probable  formula  is  A1203,  and  in 
this  case  the  atomic  weight  of  the  metal  is  Al  =  27.  (See  also 
"Double  Salts  and  Isomorphism"  in  Chap.  IX.) 

Atomic  Heat. — The  quantity  of  heat  which  a  given  weight 
of  a  metal  absorbs  while  its  temperature  rises  through  a  given 
range,  is  defined  as  the  specific  heat-  of  the  metal.  The 
quantity  of  heat  required  to  raise  the  same  weight  of  water 
through  the  same  range  of  temperature  is  taken  as  the  unit. 
Now,  it  is  found  that  the  quantity  of  heat  absorbed  by  a  solid 
metal  is  inversely  proportional  to  its  atomic  weight.  Or,  if 
weights  of  several  metals  proportional  to  their  atomic  weights 
are  taken,  the  same  quantity  of  heat  is  required  to  raise  each 
of  these  weights  through  the  same  temperature.  From  this 
it  appears  that  the  specific  heat  of  a  metal  multiplied  by  its 
atomic  weight  furnishes  a  quantity  which  is  the  same  for  all 
metals.  6-2  is  a  mean  value  for  this  quantity,  which  may  be 


EQUIVALENTS  AND  ATOMIC  WEIGHTS  129 

defined  as  the  atomic  heat  of  the  metals,  or  as  the  specific 
heat  of  their  atomic  weights.  This  is  summarized  in  the 
equation : 

Specific  heat  x  atomic  weight     =          6 -2 

6*2 

Therefore  atomic  weight  =  =. -— 

Specific  heat 

Thus,  when  the  specific  heat  of  a  metal  has  been  determined, 
this  principle,  which  is  known  as  the  law  of  Dulong  and 
Petit,  must  afford  considerable  help  in  fixing  its  atomic 
weight. 

The  solid  non-metals  are  usually  included  in  the  general 
statement,  but  for  carbon,  boron,  and  silicon  the  atomic  heat 
is  abnormal,  arid  varies  for  the  different  allotropic  modifications 
of  these  elements.  If,  however,  their  specific  heats  are  deter- 
mined at  high  temperatures,  when  their  molecules  probably 
have  a  simpler  structure,  a  much  closer  approximation  to  the 
mean  atomic  value  is  obtained.  So  that  the  principle  may 
probably  be  said  to  include  all  solid  elements. 

Chemical  Considerations.— The  formula  for  water  was 
at  first  written  HO,  and  is  correct  in  so  far  as  it  represents 
the  equivalent  weights  of  the  two  elements  in  the  compound — 
i.e.,  1  :  8.  But  do  these  symbols  represent  atoms  according 
to  the  present  view  of  the  atomic  hypothesis  1  One  of  the 
first  principles  of  the  atomic  theory  is  that  when  a  chemical 
change  takes  place  in  the  properties  of  a  substance,  it  is  the 
individual  molecules  which  undergo  the  change,  and  this  is 
true  for  both  elements  and  compounds.  In  Exp.  16  it  is 
proved  that  when  sodium  acts  upon  water,  hydrogen  gas 
is  liberated.  Where  does  this  hydrogen  come  from?  From 
the  molecules  of  the  water  is  the  only  satisfactory  explanation. 
Now,  it  is  found  that  water  disappears  in  the  same  proportion 
as  hydrogen  appears,  and,  if  a  given  weight  of  sodium  is  used, 
an  equivalent  weight  of  hydrogen  is  obtained.  Also,  if  the 
residual  liquid  is  evaporated  to  get  rid  of  the  excess  of  water, 
a  white  solid  residue  is  obtained.  This  is  the  common 

9 


130  METALLURGICAL  CHEMISTRY 

compound  known  as  caustic  soda,  which,  since  nothing  but 
hydrogen  and  excess  of  water  have  escaped,  must  contain  the 
remnants  of  the  water  molecules,  together  with  the  sodium 
used  to  bring  about  the  change.  Now,  if  this  solid  is  heated 
with  more  sodium,  a  further  quantity  of  hydrogen  is  obtained, 
and  this,  too,  must  come  from  the  remnants  of  the  water 
molecules  associated  with  the  first  quantity  of  the  metal. 
The  residue  from  this  change  is  a  solid  when  cold,  and  is 
quite  free  from  hydrogen.  From  this  it  appears  that  hydrogen 
can  be  expelled  from  the  water  molecules  in  two  distinct 
portions.  Other  metals  will  displace  this  second  proportion 
of  hydrogen  from  caustic  soda,  and  their  action  can  be  more 
readily  controlled. 

EXP.  73. — Weigh  2  grams  of  caustic  soda,  and  put  it  together 
with  20  c.c.  of  water  into  the  flask  A  (Fig.  21).  Add  0  25  gram  of 
aluminium  foil,  and  gently  heat  the  mixture  on  the  hot  gauze.  Test 
the  gas  collected  in  C.  The  characteristic  explosion  of  a  mixture  of 
air  and  hydrogen  is  obtained.  Hydrogen  is  displaced  from  the 
caustic  soda  by  the  aluminium,  and  a  double  oxide  of  aluminium 
and  sodium  is  formed. 

The  inference  to  be  drawn  from  these  experiments  is,  that 
since  hydrogen  can  be  expelled  from  the  water  molecules  in 
two  distinct  portions,  there  must  be  at  least  two  atoms  of 
hydrogen  in  each  molecule  of  water.  No  experiment  has  yet 
been  made  in  which  the  gas  is  displaced  in  more  than  two 
portions. 

Further,  it  is  possible  to  displace  oxygen  from  water,  but 
it  all  comes  out  at  once.  The  simple  explanation  of  this  is 
that  there  is  only  one  atom  of  oxygen  in  the  water  molecule. 
If,  then,  symbols  are  used  to  represent  atoms,  there  is  clearly 
only  one  way  to  write  the  formula  for  water — i.e.,  H20.  This 
is  also  borne  out  by  the  results  of  Exp.  19,  by  which  it  is 
proved  that  water  contains  hydrogen  and  oxygen  in  the 
proportion  of  2  to  1  by  volume.  For  by  Avogadro's  law  equal 
volumes  of  gases  under  normal  conditions  contain  the  same 
number  of  molecules.  Therefore,  the  water  molecules  must 


EQUIVALENTS  AND  ATOMIC  WEIGHTS  131 

supply  double  the  number  of  hydrogen  atoms  to  make  up 
the  double  volume  of  that  gas  obtained  from  the  water. 

If,  then,  the  formula  H20  is  admitted,  and  the  atomic 
weights  are  based  upon  H  =  1  for  the  standard,  it  is  evident 
that  0  =  16,  or  the  atomic  weight  of  oxygen  is  double  its 
chemical  equivalent. 

The  properties  of  water  are  so  different  from  those  of  its 
constituents  that  there  must  be  a  very  intimate  relation 
between  the  atoms  of  oxygen  and  hydrogen  in  the  molecules 
of  water.  Some  very  powerful  influence  is  at  work  holding 
the  atoms  together,  and  this  must  reside  in  the  atoms  them- 
selves, for  the  energy  of  combination  cannot  exist  apart  from 
the  constituents  of  the  compound.  It  is  not  difficult,  then,  to 
think  of  the  atoms  in  a  molecule  as  clinging  together,  merging 
their  individual  properties  into  each  other,  and  thus  producing 
a  new  particle,  with  well-defined  properties  of  its  own.  If  this 
is  so,  an  atom  of  oxygen  must  have  twice  the  clinging  power 
of  an  atom  of  hydrogen,  and  this  may  be  shown  graphically 
by  writing  the  formula  for  water  thus,  H — 0 — H.  Several 
names  have  been  given  to  this  important  property  of  atoms, 
and  the  best  are  given  below. 

Atom-combining1  Power,  Valency,  and  Atomicity.— 

When  sodium  pushes  out  part  of  the  hydrogen  from  water,  it 
evidently  displaces  an  atom  from  each  molecule  acted  upon, 
and  an  atom  of  sodium  takes  its  place.  This  is  shown  by 
writing  the  formula  of  caustic  soda  thus  :  Na — 0 — H ;  and  it 
is  seen  that  the  atom-combining  power  of  sodium  is  the  same 
as  that  of  hydrogen.  This  is  shown  again  when  the  second 
part  of  hydrogen  is  displaced  by  sodium,  and  the  formula 
becomes  Na — O — Na.  Again,  magnesium  or  zinc  pushes  the 
whole  of  the  hydrogen  out  of  water  at  one  operation,  and  it 
is  not  difficult  to  conceive  that  an  atom  of  magnesium  or  zinc 
enters  each  molecule  to  take  the  place  of  the  two  atoms  of 
hydrogen  displaced.  Hence  the  change  is  from  H — 0 — H  to 
Mg  =  0  or  Zn  =  O,  and  the  atom-combining  power  of  these 


132  METALLURGICAL  CHEMISTRY 

metals  is  equal  to  that  of  oxygen.  From  this  it  seems  that 
the  term  atom-displacing  power  also  expresses  this  property  of 
atoms. 

When  a  mixture  of  hydrogen  and  chlorine  is  fired,  rapid 
combination  takes  place  between  the  two  gases,  and  they  dis- 
appear in  exactly  equal  volumes  to  form  hydrochloric  acid 
gas,  the  total  volume  of  which  is  the  same  as  that  of  the  dis- 
appearing gases.  Now,  it  is  practically  certain  that  the 
molecules  of  the  elementary  gases  contain  two  atoms  ;  and  by 
Avogadro's  law  the  equal  volumes  contain  the  same  number 
of  molecules.  Therefore,  it  may  be  considered  that  each  pair 
of  hydrogen  and  chlorine  molecules  in  the  mixture  divide  up 
and  have  their  atoms  rearranged.  Thus  : 


| jy_H_|  +    j  Cl  |_ClJ    =    |jlJ_Cl  I   +    j  H  |  Cl  | 

2  vols.  2  vols.  2  vols.  2  vi  Is. 

This  points  to  chlorine  and  hydrogen  combining  together 
atom  to  atom.  They  have,  therefore,  the  same  combining 
power.  This  information  is  useful,  as  it  gives  another  standard 
by  which  to  measure  the  combining  power  of  metals,  since 
they  all  unite  with  chlorine,  and  the  relative  number  of  atoms 
in  the  molecules  of  their  chlorides  can  be  determined. 

Bismuth  unites  with  chlorine  and  forms  a  compound,  the 
formula  of  which  may  be  written  graphically,  thus  : 

Cl— Bi— Cl,  or  BiCl3 

.Cl 
Tin  unites  with  chlorine  to  form  the  compound  : 

Cl 

Cl — Sn — Cl,  or  SnCl4,  stannic  chloride. 


It  may  be  noticed  here  that  tin  also  unites  with  two  atoms 
of  chlorine  to  form  the  compound  SnCl2,  stannous  chloride. 
So  that  the  total  combining  power  of  an  elementary  atom 


EQUIVALENTS  AND  ATOMIC  WEIGHTS 


133 


need  not  be  always  used  up  in  the  formation  of  a  molecule. 
But  the  compound  formed  is  "  unsaturated,"  and  can  take  up 
a  second  proportion  of  the  other  element  to  form  the  higher 
compound. 

Reference  to  the  table  on  p.  125  will  now  make  it  clear  that 
the  number  which  expresses  the  ratio  of  the  chemical  equiva- 
lent of  an  element  to  its  atomic  weight,  also  gives  its  atom- 
combining  power  or  valency.  This  part  of  the  subject  should 
be  carefully  studied,  and  not  put  aside  until  the  principles  are 
quite  clear. 

Atomic  weight  and  valency  are  fundamental  properties,  and 
are  in  constant  use  for  determining  the  formulae  by  which 
compounds  are  to  be  represented.  Chap.  IX.  should  be  studied 
in  conjunction  with  this  one,  for  a  further  development  of  the 
subject  of  chemical  formulae.  The  following  table  contains 
the  common  elements  arranged  in  groups  according  to  their 
valency,  which  is  indicated  by  a  number  attached  to  the 
typical  element  of  each  group.  When  an  element  appears  in 
more  than  one  group  it  is  an  indication  that  its  active  valency 
varies  in  different  compounds. 

TABLE  OF  ATOM-COMBINING  POWER  OR  VALENCY  OF 
ELEMENTS. 


MONADS. 

DlADS. 

TRIADS. 

TETRADS. 

PENTADS.   1    HEXADS. 

H=l. 

0-2. 

O=2. 

N  =  3. 

C  =  4 

P-ft.                 S-fi 

Hydrogen 
Chlorine 
Bromine 
Iodine 

Oxygen 
Carbon 
Sulphur 
Barium 

Magnesium 
Manganese 
Mercury 
Nickel 

Nitrogen 
Boron 
Phosphorus 
Antiinonv 

Carbon          Phosphorus1  Sulphur 
Sulphur         Nitrogen       Chromium 
Silicon           Antimony     Cobalt 
Aluminium  Arsenic          Iron 

Fluorine 

Calcium         Platinum      Arsenic 

Chromium     Bismuth 

Platinum 

Potassium 

Strontium    Iron 

Bismuth 

Iron 

Tungsten 

Sodium 
Silver 

Copper          Lead 
Cadmium    i  Tin 

Gold 

Lead 
Platinum 

Cobalt           Zinc 

Tin 

Chromium 

SUMMARY. 

The  chemical  equivalents  of  the  elements  in  general,  or 
their   quantitative   relations   to    each    other,  form   the   solid 


134  METALLUKGICAL  CHEMISTKY 

foundation  upon  which  all  systematic  practical  work  is  based. 
Mere  observation  is  useful  for  indicating  the  way,  but  it  is 
only  when  one  begins  to  weigh  and  measure  that  the  great 
possibilities  of  the  subject  become  apparent.  The  selection  of 
hydrogen  as  the  standard  weight  renders  comparison  easy. 
The  chemical  equivalent  of  an  element  is  purely  the  result  of 
experiment,  and  is  quite  independent  of  any  theory  of  the 
constitution  of  matter.  The  atomic  weight  of  an  element, 
on  the  other  hand,  assumes  the  existence  of  atoms,  and  is 
dependent  on  the  atomic  theory  for  its  definition ;  but  it  is 
none  the  less  certain  on  that  account,  for  it  is  simply  the 
chemical  equivalent  multiplied  by  a  whole  number.  The 
selection  of  this  number  for  a  particular  element  is  also  based 
upon  experimental  methods,  and  the  number  itself  indicates 
the  atom-combining  power  or  valency  of  the  element.  The 
great  care  which  has  been  exercised  in  the  determination  of 
atomic  weights  renders  them  perfectly  reliable  for  use  in  the 
most  important  calculations,  and  every  confidence  may  be 
placed  in  them. 

QUESTIONS. 

1.  What  do  you  understand  by  the  chemical  equivalent  of 
an  element  ? 

2.  Describe  two  methods  by  which  the  chemical  equivalent 
of  copper  can  be  determined. 

3.  What  is  meant  by  the  valency  of   an  element?     Give 
examples. 

4.  Describe  experiments  in  which  hydrogen  is  displaced  from 
water  in  two  distinct  operations. 

5.  What  is  the  relation  between  the  chemical  equivalent  of 
an  element  and  its  atomic  weight  ? 

6.  Explain  the  difference  between  the  atom-combining  power 
of  an  element  and  its  combining  weight. 


CHAPTER  IX 
OXIDES,  ACIDS,  AND  SALTS 

ALL  the  metals  combine,  either  directly  or  indirectly,  with 
oxygen  to  form  oxides;  but  those  oxides  which  cannot  be 
formed  by  direct  oxidation  are  unstable  bodies  of  little  practical 
importance.  The  common  metals  which  do  not  oxidize,  when 
exposed  to  air  or  oxygen  under  suitable  conditions,  are  silver, 
gold,  and  platinum. 

A  few  of  the  common  oxides  unite  directly  with  water  to 

form  compounds,  called  hydroxides  or  hydrates.    These 

definite  bodies  contain  the  elements  of  the  oxides  and  of 
water.  Thus,  sodium  oxide,  Na20,  and  potassium  oxide, 
K20,  form  very  stable  compounds  with  water,  which  are  not 
decomposed  at  a  bright  red  heat.  This  kind  of  change  is 
readily  shown  : 

Na20  +  H20  =  2NaHO. 

Sodium 
hydroxide 

Calcium  oxide  (quicklime),  GaO,  and  the  corresponding  oxides 
of  barium  and  strontium,  combine  readily  with  water,  and  a 
considerable  amount  of  heat  is  developed  during  the  reaction. 
Thus: 

CaO  +  H20=Ca(HO)2. 

Calcium 
hydroxide 

These  hydroxides,  however,  are  decomposed  at  a  red  heat,  and 
the  water  escapes,  leaving  the  original  oxides  as  a  residue. 
The  change  shown  in  the  last  equation  is  then  reversed.  The 
evolution  of  heat  is  readily  observed  on  moistening  a  piece  of 
quicklime  with  water. 

135 


136  METALLURGICAL  CHEMISTRY 

Hydroxides  which  can  be  formed  by  the  direct  union  of 
oxides  with  water  are  more  or  less  soluble  in  the  liquid,  and 
their  solutions  have  very  characteristic  properties.  They 
produce  a  soapy  feeling  when  rubbed  between  the  thumb  and 
finger,  and  the  very  soluble  ones  have  a  caustic  or  burning 
action  upon  the  skin.  They  all  turn  reddened  litmus  blue. 
This  is  called  an  alkaline  reaction,  and  the  compounds  pro- 
ducing it  are  alkaline  hydroxides. 

The  soapy  feeling  mentioned  above  is  really  due  to  the 
formation  of  a  soap,  by  a  reaction  between  the  alkali  and  the 
fatty  matter  in  the  skin.  The  common  alkalies,  caustic  soda 
and  caustic  potash,  are  largely  used,  on  account  of  this  property, 
for  the  removal  of  grease  from  various  articles.  The  soap 
formed  is  readily  washed  away,  and  the  grease  thus  removed. 
Hydroxides  of  metals,  the  oxides  of  which  do  not  combine 
directly  with  water,  can  be  formed  in  an  indirect  way.  They 
are  insoluble  in  water,  and  when  a  solution  of  a  soluble 
hydroxide  is  added  to  a  solution  of  a  salt  of  one  of  these 
metals,  the  hydroxide  of  the  metal  is  generally  formed,  and 
precipitated  from  the  solution. 

EXP.  74. — Put  a  crystal  of  copper  sulphate  into  a  test-tube,  and 
dissolve  it  in  a  little  water ;  add  a  solution  of  caustic  soda,  a  little 
at  a  time,  as  long  as  a  precipitate  forms  ;  shake  well,  and  allow  the 
precipitate  to  settle.  Repeat  the  experiment  with  a  solution  of 
'ferric  chloride  in  place  of  the  copper  sulphate. 

The  green  and  brownish-red  precipitates  are  hydroxides, 
and  the  changes  are  shown  by  the  equations : 

CuS04  +  2NaHO  =  Cu(HO)2  +  Na,S04. 
Fe2Cl6  +  GNaHO  «  Fe2(HO)6+  GNaCl. 

Basic  Oxides. — Metallic  oxides  which  form  hydroxides, 
either  soluble  or  insoluble,  corresponding  to  the  above,  are 
called  basic  Oxides,  and  it  is  very  probable  that,  if  the  in- 
soluble ones  could  be  dissolved  in  water,  their  solutions  would 
give  an  alkaline  reaction.  The  following  are  examples  of  basic 
oxides  : 

K20,  CaO,  MgO,  HgO,  PbO,  ZnO,  A1203,  Fe203. 


OXIDES,  ACIDS,  AND  SALTS  137 

Peroxides. — But  some  metals  form  more  than  one  oxide, 
and  it  is  then  necessary  to  enquire  if  all  the  oxides  of  a  parti- 
cular metal  are  basic  in  character. 

Compare  PbO  and  Pb02 ;  BaO  and  Ba02 ;  MnO  and  Mn02. 
The  second  oxide  in  each  pair  gives  off  oxygen  gas  when  it  is 
heated  alone,  or  with  dilute  sulphuric  acid ;  and  when  heated 
with  hydrochloric  acid  causes  the  liberation  of  chlorine  from 
the  acid.  For  example  : 

2Ba02  =  2BaO  +  02. 
Mn02  +  4HC1  =  MnCl2  +  2H20  +  C12. 

These  higher  oxides  are  not  true  basic  compounds,  and  are 
usually  called  peroxides. 

Complex  Oxides. — Another  class  of  oxides,  of  which  red 
lead,  Pb3O4,  and  magnetic  oxide  of  iron,  Fe304,  are  examples, 
are  sometimes  regarded  as  being  a  combination  of  two  oxides 
of  the  same  metal.  Compare  Pb304  and  2PbO.Pb02 ;  Fe3O4 
and  FeO.Fe00Q.  They  are  not  true  basic  oxides. 

Z      o  «/ 

Acid-forming"  Oxides. — A  few  metallic  oxides  are  even 
more  highly  oxidized  than  the  peroxides.  A  well-known 
example  of  such  compounds  is  the  highest  oxide  of  chromium, 
Cr03.  This  oxide  when  in  solution  has  distinctly  acid  pro- 
perties, and  is  altogether  different  from  the  basic  oxides  of  the 
same  metal.  It  is  an  acid-forming  oxide. 

Metallic  oxides  may,  then,  be  divided  into  four  classes  : 
(1)  basic  oxides;  (2)  peroxides;  (3)  complex  oxides;  (4)  acid- 
forming  oxides. 

Oxides  Of  Non-Metals. — The  most  powerful  acid-forming 
oxides  are  to  be  found  among  the  oxides  of  the  non-metals. 
They  unite  with  water  to  form  hydroxides,  which  are  mostly 
soluble  in  water,  and  form  solutions  having  very  characteristic 
properties.  The  solution  of  an  acid  hydroxide  has  a  SOUP 
taste,  and  turns  litmus  *  red.  This  is  the  characteristic  acid 

*  Litmus  is  a  vegetable  colouring  matter.  It  dissolves  readily  in  water, 
and  either  the  solution  itself  or  porous  paper  stained  with  it  may  be  used 
for  testing:. 


138  METALLUEGICAL  CHEMISTKY 

reaction,  as  distinguished  from  the  alkaline  or  basic  reaction, 
in  which  reddened  litmus  is  turned  blue.  The  formation  of 
an  acid  hydroxide  is  represented  thus  : 


24. 

Sulphur  Sulphuric 

trioxide  acid 

A  few  of  the  non-metals  form  peroxides,  and  some  of  the  non- 
metallic  oxides  are  neutral  in  character. 

Nomenclature  Of  Oxides.  —  When  there  is  only  one  oxide 
of  a  particular  element  no  difficulty  is  experienced  in  giving 
it  a  name,  as  MgO,  magnesium  oxide  ;  ZnO,  zinc  oxide.  But, 
when  there  are  two  or  more  oxides  of  the  same  element,  dis- 
tinguishing names  must  be  used.  Thus,  copper  forms  two 
compounds,  commonly  called  the  red  and  black  oxides,  from 
their  respective  colours.  The  systematic  names  are  cuprOUS 
oxide,  Cu2O,  and  cuprie  oxide,  CuO.  Similarly,  we  have 
ferrous  and  ferric  oxides,  FeO  and  Fe203.  The  termination 
-OUS  indicates  the  oxide  containing  the  smaller  proportion, 
and  the  termination  -ic  the  larger  proportion,  of  oxygen.  The 
prefix  sesqui-  is  also  used  to  denote  the  ratio  2  to  3.  Sesqui- 
oxide  of  iron  is  ferric  oxide,  Fe203. 

The  same  terminations  are  used  to  distinguish  hydroxides  of 
the  same  element  from  each  other.  There  are  ferrous  and 
ferric  hydroxides;  and  the  two  important  acid  hydroxides 
of  sulphur  are  sulphurous  and  sulphuric  acids,  H2S03  and 
H2S04. 

Another  useful  way  of  naming  is  to  put  a  numerical  prefix 
before  the  word  "  oxide,"  and  also  before  the  name  of  the  other 
element,  if  necessary.  The  prefixes  indicate  the  number  of 
atoms  of  oxygen  or  of  the  other  element  in  the  formula. 

Thus  sulphur  dioxide,  S02  ;  sulphur  trioxide,  S03  ;  and 
triferric  tetroxide,  Fe304.  Other  examples  are  carbon  mon- 
oxide, CO  ;  carbon  dioxide,  C02  ;  chromium  trioxide,  Cr03  ; 
nitrogen  tetroxide,  N2O4  ;  phosphorus  pentoxide,  P205. 

Acid-forming  oxides  are  often  called  anhydrides,  as 
sulphuric  anhydride,  S03. 


OXIDES,  ACIDS,  AND  SALTS  139 

Many  compounds  have  common  names,  with  which  the 
student  will  become  familiar,  and,  with  a  little  practice,  he 
will  have  no  difficulty  in  giving  the  systematic  name  to  any 
compound,  when  he  knows  its  formula. 

Basic  and  acid-forming  oxides  are  the  most  important,  and 
it  may  be  stated  definitely  with  regard  to  their  distinctive 
properties  that,  if  there  are  two  basic  oxides  of  the  same 
metal,  the  lower  oxide  is  the  stronger  base.  Thus,  FeO  is 
a  stronger  base  than  Fe203.  On  the  other  hand,  of  two  acid- 
forming  oxides  of  the  same  element  the  higher  oxide  forms 
the  Stronger  acid.  Thus,  sulphuric  acid  is  more  powerful 
than  sulphurous  acid.  With  some  of  the  metals,  both  acid- 
forming  and  basic  oxides  are  obtained.  This  is  so  with  the 
oxides  of  chromium  : 

Chromous  oxide,  CrO ;  chromic  oxide,  Cr203; 

Strongly  basic  Basic 

chromium  trioxide,  CrO3. 

Acid -forming 


REACTIONS  BETWEEN  BASIC  OXIDES  AND  ACIDS. 

The  acid  hydroxides  are  such  well-defined  compounds  that 
they  were,  for  a  long  time,  considered  to  be  the  typical  acids, 
and  the  oxygen  they  contain  absolutely  essential  to  the  com- 
position of  an  acid.  But  when  it  was  conclusively  proved 
that  hydrochloric  acid  contains  no  oxygen,  this  view  had  to 
be  modified.  Now,  clear  proof  has  been  furnished  that,  when 
some  of  the  metals  are  presented  to  the  common  acids,  they 
displace  hydrogen  and  take  its  place,  forming  salts  of  the 
acids.  So  that  it  would  appear  that  the  essential  element  in 
an  acid  is  the  hydrogen  which  can  be  displaced  from  it  by  a 
metal.  Now,  only  a  comparatively  small  number  of  metals 
will  displace  hydrogen  from  acids  in  general,  but  a  large  number 
of  basic  oxides  and  hydroxides  will  react  with  most  acids  and 
form  salts  of  the  acids  used.  Thus  copper  does  not  dissolve 
in  either  dilute  sulphuric  or  hydrochloric  acid,  but  cupric 


140  METALLURGICAL  CHEMISTRY 

oxide  does  dissolve,  and  forms  the  copper  salts  of  the  acids. 
Thus  : 

1  .  CuO  +  2HC1  =  CuCl2  +  H0O. 

2. 


It  is  seen  from  the  equations  that  the  hydrogen  of  the  acid  is 
replaced  by  the  metal  of  the  oxide,  and  that  the  displaced 
hydrogen  unites  with  the  oxygen  of  the  oxide  to  form  water. 
This  is  the  general  reaction. 

The  experiments  which  follow  are  intended  to  make  the 
beginner  practically  acquainted  with  the  commoner  reactions 
between  oxides  and  acids. 

EXP.  75.  —  Weigh  roughly  3  grams  of  caustic  potash,  put  it  into 
a  porcelain  basin  with  10  c.c.  of  water,  and  stir  the  liquid  with  a 
glass  rod  until  the  solid  is  dissolved.  Test  the  solution  with  apiece 
of  red  litmus-paper,  and  leave  the  paper  sticking  to  the  side  of  the 
dish,  so  that  a  slight  movement  will  bring  the  liquid  into  contact 
with  it.  Now  add  dilute  nitric  acid  (1  to  1)  a  little  at  a  time,  stirring 
with  a  glass  rod,  until  the  blue  colour  of  the  litmus-paper  is  just 
turned  red  when  the  solution  is  brought  into  contact  with  it. 
Towards  the  end  the  acid  must  be  added  drop  by  drop.  The  first 
portion  of  the  acid  may  be  run  in  from  a  test-tube,  and  the  remainder 
from  a  piece  of  glass  tube  pulled  off  to  a  jet.  The  dropping-tube  is 
put  into  the  test-tube  containing  the  acid,  and  when  some  of  the 
liquid  has  entered,  the  top  of  the  tube  is  closed  by  the  finger,  and 
the  tube  itself  lifted  out.  By  moving  the  finger  the  acid  is  allowed 
to  escape  drop  by  drop.  When  the  action  is  finished,  remove  the 
litmus-paper,  evaporate  the  solution  down  to  about  15  c.c.,  and  set 
it  aside  to  crystallize.  Examine  the  crystals,  and  compare  them 
with  the  saltpetre  crystals  in  the  laboratory  bottle. 

In  this  experiment  it  is  clear  that  the  basic  properties  of 
the  alkali  and  the  acid  properties  of  the  acid  have  entirely 
disappeared,  along  with  the  bodies  themselves,  and  a  new  body 
has  been  formed.  The  reaction  is  expressed  thus  : 

KHO  +  HN03  =  KN03  +  H20. 

Basic  Acid  Neutral          Water 

hydroxide     hydroxide  salt 

This  equation  not  only  shows  the  general  change,  but  also 
expresses  these  quantitative  relations  between   the   reacting 


OXIDES,  ACIDS,  AND  SALTS  141 

compounds.  These  relations  are  just  as  definite  as  those  which 
exist  between  the  elements  in  an  oxide,  or  between  the  amount 
of  a  dissolved  metal  and  the  gas  liberated  during  its  dissolution. 
This  is  easily  proved  by  using  solutions  of  acids  and  alkalies 
containing  known  weights  of  the  reacting  compounds,  and 
finding  the  volumes  of  the  liquids  required  to  be  added  together 
to  bring  about  neutralization. 

Solutions  containing  10  grams  per  litre  of  each  of  the 
following  pure  compounds  are  easily  prepared :  sulphuric, 
nitric,  and  hydrochloric  acids,  caustic  soda  and  caustic 
potash.  Directions  for  the  preparation  of  one  of  them  is 
given  on  p.  143,  but  they  should  be  ready  prepared  for  the 
student's  use. 

The  following  exercises  may  be  made  successful  with 
moderate  care,  and  will  give  a  good  insight  into  the  quantita- 
tive relations  which  exist  during  neutralization.  Litmus  has 
been  tised  so  far  as  an  indicator  of  the  acid  or  alkaline  character 
of  a  solution  ;  but,  for  the  work  now  to  be  described,  methyl 
Orange*  is  to  be  preferred,  as  it  is  not  affected  by  carbon 
dioxide  in  solution.  It  is  turned  greenish-yellow  by  an  alkali, 
but  changes  to  a  bright  red  on  the  addition  of  a  very  slight 
excess  of  an  acid. 

EXP.  76.— Wash  out  the  burette  (Ch.  xv.)  with  a  little  of  the 
standard  acid  solution,  fix  it  in  the  clip,  and  fill  it  up  to  the  mark 
with  the  same  solution.  Transfer  10  c.c.  of  the  caustic  soda  solution 
to  a  conical  flask,  by  means  of  a  10  c.c.  pipette,  add  about  100  c.c. 
of  water  to  the  flask,  and  one  or  two  drops  of  methyl  orange.  Note 
the  level  of  the  acid  solution  in  the  burette,  and  then  run  a  little  at 
a  time  into  the  flask,  shaking  it  between  each  addition,  until  the 
yellow-green  colour  of  the  solution  changes  to  red.  Care  must  be 
exercised  towards  the  end  of  the  reaction,  as  one  drop  of  the  acid 
solution  is  sufficient  to  change  the  colour.  Bead  off  the  volume  of 
acid  solution  added.  Eepeat  the  experiment  with  a  second  10  c.c. 
of  the  soda  solution,  and  compare  the  two  results. 

An  exactly  similar  experiment  may  then  be  made  with 
caustic  potash  solution. 

*  The  compound  known  as  Porrier's  Orange  No.  3  is  the  best ;  1  gram 
per  litre  makes  a  good  solution. 


142  METALLURGICAL  CHEMISTRY 

EXAMPLE. — 10  c.c.  of  caustic  soda  solution  required  12 '3  c.c.  of 
sulphuric  acid  solution,  and  10  c.c.  of  caustic  potash  solution 
required  8'7  c.c.  of  sulphuric  acid  solution. 

Since  each  solution  is  known  to  contain  10  grams  of  the 
compound  in  1,000  c.c ,  O'Ol  gram  is  contained  in  1  c.c. 

.'.01  gram  of  caustic  soda  neutralizes  0  123  gram  of  sul- 
phuric acid,  and  0-1  gram  of  caustic  potash  0087  of  the 
acid. 

These  weights  may  be  compared  directly  with  the 
equations  : 

2NaHO  +  H2S04  =  Na2S04  +  2H20. 

Sodium 
sulphate 

2KHO  +  H2S04  =  K2S04+  2H2O. 

Potassium 
sulphate 

Q.I  03 

Also  ,r^r—-  =  1*41  =  the  ratio  of  the  weights  of  acid  required 
0*087 

to  neutralize  the  same  weight  of  the  alkalies. 

Experiments  with  nitric  acid  and  with  hydrochloric  acid 
may  be  made  in  the  same  way,  and  the  results  compared. 

EXAMPLE. — 10  c.c.  of  caustic  soda  solution  required  15'6  c.c.  of 
nitric  acid  solution,  and  10  c.c.  of  caustic  potash  solution  required 
11 '2  c.c.  of  nitric  acid  solution. 

NaHO  +  HN03  =  NaNo3  +  H20. 

Sodium 
nitrate 

Also B  =1-39  =  ratio  of  weights  of  nitric  acid  required  to 

neutralize  the  same  weight  of  the  alkalies. 

EXAMPLE. — 10  c.c.  of  the  caustic  soda  solution  required  9'1  c.c. 
of  hydrochloric  acid  solution,  and  10  c.c.  of  the  caustic  potash 
solution  required  6'5  c.c.  of  hydrochloric  acid  solution. 

KHO  +  HC1  =  KC1  +  H20. 

Potassium 
chloride 

A.Q1 

The  ratio  is    4±-l-41. 


OXIDES,  ACIDS,  AND  SALTS  143 

It  is  evident  from  the  above  results  that,  if  equal  weights 
of  two  bases  are  used  to  neutralize  different  acids,  the  ratio 
of  the  weights  of  the  acids  neutralized  is  constant.  This  fact 
was  recognised  as  early  as  the  latter  part  of  the  eighteenth 
century,  and  work  done  in  this  direction  by  Kichter  and  others 
helped  to  lay  the  foundation  of  the  first  law  of  chemical  com- 
bination, which  is  stated  on  p.  47. 

Preparation  of  Solutions. — In  making  an  acid  or  an 
alkaline  solution  of  known  strength,  it  is  convenient  to  add 
more  than  the  required  quantity  of  the  compound,  and  after 
determining  the  actual  weight  present  in  a  given  volume,  to 
bring  the  whole  to  the  proper  strength  by  the  addition  of  a 
calculated  volume  of  water.  The  following  proportions  may 
be  used,  and  a  large  or  small  volume  of  the  solution  prepared 
as  required. 

Concentrated  sulphuric  acid          7  c.c. 
nitric  acid  12   , 


,     .      , ,     .       . ,  __  ,  800  c.c.  of  water 

hydrochloric  acid  25   ,,        v  , 

in  each  case. 
Caustic  soda  10  grams 

„       potash  10      „     J 

The  strength  of  an  acid  solution  is  usually  determined  by 
taking  advantage  of  its  reaction  with  pure  sodium  carbonate. 
With  sulphuric  acid  this  is  : 

Na2C03     +     H2S04 
46  +  12  +  48     2  +  32  +  64 


106  98 

The  carbonate  during  its  decomposition  neutralizes  the  acid, 
and  by  using  the  methyl  orange  indicator,  the  exact  amount 
of  acid  to  be  added  is  readily  controlled.  From  the  equation 
it  is  clear  that  106  parts  by  weight  of  Na2C03  neutralize  98 

1  Ofi 

parts   by  weight  of   H2S04;    .*.  ——=1-082  =  the  weight   of 

t/o 

carbonate  required  to  neutralize  1  part  by  weight  of  the  acid ; 
and  0'2164  gram  is  required  for  0'2  gram  of  the  acid. 


144  METALLUKGICAL  CHEMISTEY 

If,  then,  0-2164  gram  of  pure  sodium  carbonate  is  dissolved 
in  about  100  c.c.  of  water,  one  drop  of  methyl  orange  added, 
and  the  sulphuric  acid  solution  run  in  from  a  burette  until 
the  colour  change  takes  place,  the  volume  of  the  acid  solution 
used  will  contain  exactly  0'2  gram  of  H2S04.  The  whole 
volume  of  the  remaining  solution  may  then  be  measured,  and 
the  necessary  volume  of  water  added  to  bring  it  up  to  the 
proper  strength. 

If  V  =  the  volume  of  the  remaining  acid  solution,  and  v  = 
the  volume  used  to  neutralize  0'2164  gram  of  carbonate, 

y 

then  — (20 — v)  =  volume  of  water  to  be  added. 

EXAMPLE. — Volume  of  remaining  solution  =  750  c.c. ;  volume 
used  to  neutralize =17  c.c. 

/.  71y(20--17)  =  75!j^  =  132  c.c.  of  water  to  be  added. 

The  strengths  of  the  other  acid  solutions  can  be  determined 
in  a  similar  manner  by  using  the  equations : 

:Na2C03  +  2HN03  =  2NaN03  +  C02  +  H20. 
Na2C03  +  2HC1  =  2NaCl  +  C02  +  H20. 

The  atomic  weights  may  be  taken  from  the  table  in  the 
appendix.  The  strength  of  the  alkaline  solutions  are  then 
readily  determined  with  one  of  the  acid  solutions.  Twenty  c.c. 
of  the  soda  solution  are  transferred  to  a  flask,  diluted  with 
water,  and  one  drop  of  methyl  orange  added.  The  standard 
acid  is  then  run  in  until  the  colour  changes.  For  hydro- 
chloric acid  and  caustic  soda  the  reaction  is  expressed  by  the 
equation  : 

NaHO     +     HC1     =     NaCl  +  H20. 
23  +  1  +  16       1+35-5 

~~40~~          ^3(^T 
1  gram  of  HC1=  ^  =  1-096  gram  of  NaHO. 


OXIDES,  ACIDS,  AND  SALTS  145 

Then  the  number  of  cubic  centimetres  of  acid  required  multi- 
plied by  0-01096  gives  the  weight  of  NaHO  in  20  c.c.  of  the 
solution,  and  a  simple  calculation  determines  the  volume  of 
water  to  be  added  to  the  alkaline  solution  to  bring  it  down  to 
1  c.c.  =0-01  gram  of  NaHO.  Since  each  solution  contains 
10  grams  of  the  pure  compounds  in  1,000  c.c.,  the  bottles 
may  be  labelled  as  follows  : 

Sulphuric  acid,  1  c.c.  =  0'01  gram  H2S04,  etc. 

NEUTRALIZATION  BY  INSOLUBLE  OXIDES. 

The  majority  of  the  oxides  insoluble  in  water  are  dissolved 
more  or  less  readily  by  the  common  acids,  and  the  correspond- 
ing salts  are  formed.  If,  however,  the  salt  to  be  formed  is 
insoluble  in  the  acid  solution,  dissolution  does  not  take  place, 
and  the  action  between  the  oxide  and  the  acid  is  somewhat 
slow,  but,  as  a  rule,  the  salt  is  formed  on  continued  digestion 
with  the  acid. 

Some  oxides  which  will  withstand  a  high  temperature  with- 
out fusing  dissolve  very  slowly  after  they  have  been  strongly 
ignited.  This  is  probably  due  to  a  closer  aggregation  of  the 
molecules,  by  which  they  become  more  complex  in  character 
and  more  difficult  to  decompose.  This  increase  in  molecular 
complexity  is  known  as  polymerization.  Strongly  ignited  ferric 
and  chromic  oxides  are  examples. 

Magnesia — EXP.  77.— Put  20  c.c.  of  dilute  hydrochloric  acid 
(1  to  1)  into  a  porcelain  basin,  and  test  it  with  blue  litmus-paper. 
Then  add  magnesia  a  little  at  a  time ;  stir  well  with  a  glass  rod 
between  each  addition,  and  continue  the  addition  until  some  of  the 
powder  is  left  undissolved.  Test  the  solution  with  the  litmus-paper, 
and,  if  it  turns  red,  boil  the  solution  slowly  for  a  minute  or  two  to 
see  if  the  residue  is  dissolved.  Filter  the  liquid  to  clear  it,  evaporate 
it  down  to  half  the  bulk,  and  set  it  aside  to  crystallize. 

The  crystals  are  usually  small  and  not  very  distinct.  If  the 
solution  is  evaporated  carefully  to  dryness,  and  the  dry  mass 
exposed  to  the  air,  it  absorbs  moisture,  and  becomes  quite 
moist.  It  is  a  deliquescent  salt.  The  magnesium  oxide,  MgO, 

10 


146  METALLURGICAL  CHEMISTRY 

dissolves  in  the  acid  liquid,  and  during  its  dissolution  neutral- 
izes the  acid.  It,  therefore,  establishes  its  claim  to  be  a  basic 

oxide. 

MgO  +  2HC1  P.  MgCl2  +  H20. 

Magnesium 
chloride 

A  similar  experiment  may  be  made  with  sulphuric  acid  and 
magnesia,  and  the  magnesium  sulphate  formed  is  readily 
crystallized.  Its  formula  is  MgS04,7H20. 

Copper  Oxide—  EXP.  78.  —  Put  2  grams  of  black  oxide  of  copper 
into  a  large  test-tube,  and  add  15  c.c.  of  dilute  sulphuric  acid  (1  to  6). 
Boil  the  solution  gently  until  the  whole  of  the  oxide  has  dissolved 
or  until  the  residue  becomes  distinctly  red.  Commercial  black 
oxide  of  copper  usually  contains  a  little  of  the  red  oxide,  which  is 
not  completely  dissolved  by  sulphuric  acid,  and  a  residue  of  red 
metallic  copper  is  left.  Filter  the  solution  into  a  porcelain  dish, 
and,  if  it  is  deep  blue  in  colour,  set  it;  aside  to  crystallize  ;  but  if 
only  light  blue,  evaporate  it  to  about  half  the  bulk  before  setting  it 
aside. 

CuO  +  H2S04  =  CuS04  +  H20. 

The  change  with  the  red  oxide  is  shown  thus  : 
Cu20  +  H2S04  =  CuS04  +  H20  +  Cu. 

The  blue  crystals  (blue  vitriol)  have  the  composition  CuS04, 
5H20. 

Zinc  Oxide.  —  EXP.  79.  —  Repeat  the  last  experiment,  using  zinc 
white,  ZnO,  in  place  of  black  oxide  of  copper.  Also,  dissolve  2  grams 
of  zinc  in  dilute  sulphuric  acid,  and  set  the  two  solutions  aside  to 
crystallize. 

On  comparison  the  crystals  are  found  to  be  exactly  the  same. 
Zinc  sulphate  (white  vitriol),  ZnS04,7H20,  is  formed  in  each 
case.  Compare  the  equations  : 

ZnO  +  H2S04  =  ZnS04  +  H2O. 


Oxide  of  Lead.  —  EXP.  80.  —  Put  5  grams  of  lead  oxide  into  a 
test-tube,  add  20  c.c.  of  dilute  nitric  acid  (1  to  5),  and  heat  the 
tube  gently  until  the  oxide  is  dissolved.  Filter  the  solution  into 
an  evaporating  basin,  and  set  it  aside  to  crystallize.  Examine  the 
crystals. 

PbO  +  2HNO,  =  Pb(N03)2  +  H20. 

Lead  nitrate 


OXIDES,  ACIDS,  AND  SALTS  147 

The  lead  nitrate  crystals  contain  no  water  of  crystallization, 
and  the  formula  of  the  crystalline  salt  is  Pb(N03)2.  Lead 
sulphate,  PbS04,  is  insoluble  in  dilute  sulphuric  acid,  but  on 
boiling  the  oxide  for  a  few  minutes  with  the  moderately  strong 
acid,  an  appreciable  quantity  of  the  sulphate  is  formed.  Hot 
hydrochloric  acid  also  dissolves  lead  oxide  with  formation  of 
the  chloride,  but  the  salt  largely  separates  from  its  solution  on 
cooling. 

EXP.  81. — Boil  a  little  lead  oxide  with  moderately  strong  sulphuric 
acid  (1  to  3)  for  a  few  minutes  ;  allow  the  test-tube  to  stand  until 
the  liquid  becomes  clear ;  note  the  layer  of  white  sulphate  upon  the 
yellow  residue  of  unchanged  oxide.  Repeat  the  experiment  with 
dilute  hydrochloric  acid  (1  to  1).  When  the  lead  oxide  is  all  dis- 
solved set  the  solution  aside  to  cool,  and  note  the  rapid  separation 
of  lead  chloride  crystals. 

Water  of  Crystallization. — Potassium,  sodium,  and 

lead  nitrates  crystallize  from  their  solutions  just  as  they  are 
formed ;  but  zinc,  copper,  and  magnesium  sulphates,  when 
they  crystallize,  carry  out  with  them  definite  quantities  of 
water.  This  is  called  water  of  crystallization,  and  is  associated 
with  the  salts  in  a  very  definite  manner.  It  may  be  taken 
as  a  general  statement  that  when  a  given  salt  crystallizes 
under  normal  conditions  it  always  takes  out  with  it  the  same 
proportion  of  water.  The  composition  of  a  crystalline  salt  is 
constant,  but  the  relation  between  the  salt  and  its  water  of 
crystallization  is  not  so  intimate  as  that  between  the  elements 
of  the  salt  itself.  It  is  not  so  intimate  even  as  that  between 
an  oxide  and  the  elements  of  water  in  its  hydroxide.  This  is 
shown  in  the  formula  by  keeping  the  water  separate.  Thus  : 
crystallized  copper  sulphate  is  CuS04,5H20. 

In  general,  crystalline  salts  are  freely  soluble  in  water,  but 
there  is  a  limit,  differing  with  different  salts,  to  the  quantity 
of  the  solid  dissolved  by  a  given  weight  of  water.  The  solu- 
bility of  a  salt,  or  other  solid,  is  usually  stated  in  terms  of 
parts  by  weight  of  cold  water  required  to  dissolve  1  part  by 
weight  of  the  solid.  Thus  potassium  nitrate  dissolves  in 

10—2 


148  METALLURGICAL  CHEMISTRY 

7  parts,  zinc  sulphate  in  2 -5  parts,  and  potassium  chlorate 
in  20  parts  of  cold  water. 

In  most  cases  the  salt  is  much  more  soluble  in  hot  than  in 
cold  water,  and  the  cold  saturation  point  is  thus  easily  ex- 
ceeded. As  the  solution  cools  the  excess  above  that  which 
the  cold  solution  will  hold  tends  to  separate  in  the  solid 
state — i.e.,  to  crystallize.  But  it  is  possible  for  the  solutions 
to  become  quite  cold  without  crystallization  setting  in.  Such 
a  solution  is  super-saturated,  and  will  crystallize  rapidly 
when  once  the  action  is  started.  A  small  crystal  of  the  same 
substance  dropped  in  is  usually  sufficient  to  set  up  the  action, 
and  even  a  speck  of  dust  may  form  a  nucleus  for  the  initial 
crystals  to  form  round.  Generally,  if  the  solution  is  above 
the  cold  saturation  point,  and  cools  slowly,  the  crystallization 
commences  slowly,  and  the  final  crystals  are  larger  and  better 
developed.  If  the  solution  is  too  strong  the  crystallization 
will  be  rapid  and  the  crystals  very  small,  forming  a  powder 
when  dry.  When  once  crystals  are  formed  in  the  solution 
they  continue  to  grow,  and  the  dissolved  solid  is  reduced 
somewhat  below  the  point  of  cold  saturation  ;  but  it  is  im- 
possible to  separate  the  whole  of  the  solid  from  the  solution, 
for  there  is  a  limit  at  which  the  solid  dissolves  from  the 
crystals  already  formed  as  fast  as  it  is  deposited  upon  them. 
The  residual  solution  from  which  the  crystals  have  separated 
is  called  the  "  mother  liquor." 

It  will  be  readily  understood  that  if  two  solids  differ  in 
solubility,  it  is  possible  to  separate  a  mixture  of  them  by 
repeated  crystallization,  as  the  more  soluble  one  will  not 
crystallize  as  rapidly  as  the  less  soluble  one  under  the  same 
conditions.  Also,  a  small  quantity  of  one  soluble  body  can 
be  separated  from  a  large  quantity  of  another.  These  facts 
are  taken  advantage  of  in  the  purification  of  salts  by  repeated 
crystallization.  Some  salts  dissolve  almost  as  readily  in  cold 
as  in  hot  water.  That  is,  their  solubility  increases  very  little 
with  a  rise  in  the  temperature  at  which  the  solution  is 
effected.  Common  salt  is  an  example.  Others,  again,  seem  to 


OXIDES,  ACIDS,  AND  SALTS  149 

become  less  soluble  after  a  certain  temperature  is  passed. 
This  is  so  with  sodium  sulphate,  which,  below  33°  C.,  crystallizes 
with  10  molecules  of  water,  thus  :  Na2S04,10H20.  But  when  a 
solution  which  would  barely  crystallize  at  ordinary  temperatures 
is  raised  above  33°  C.,  small  crystals  of  the  anhydrous  (water- 
less) salt,  Na2SO4,  are  deposited.  From  this  it  would  seem 
that  above  33°  C.  the  ordinary  crystallized  salt  does  not  exist, 
and  when  the  point  of  saturation  for  Na2S04  is  passed  crystals 
are  formed. 

But  this  is  really  a  case  of  equilibrium  between  the  solid  and 
the  solvent  liquid,  and  comes  within  the  scope  of  the  phase 
rule  enunciated  by  Willard  Gibbs.  For  information  on  this 
subject  the  student  is  referred  to  recent  works  on  physical 
chemistry. 

Crystals. — A  fully-developed  crystal  is  a  solid  of  definite 
geometrical  shape,  bounded  generally  by  plane  surfaces,  which 
join  each  other  and  form  solid  angles.  These  surfaces  are  the 
faces  of  the  crystals,  and  an  angle  formed  by  the  intersection 
of  two  or  more  of  these  faces  is  invariable  for  the  same 
crystalline  form.  It  is  possible  for  a  number  of  different 
substances  to  crystallize  in  the  same  form,  so  that  a  classifica- 
tion of  crystals  can  be  made.  Thus  common  salt  crystallizes 
in  the  form  of  cubes,  in  which  all  the  angles  are  right  angles, 
and  therefore  equal.  Iceland  spar  crystallizes  in  rhombs,  in 
which  the  opposite  angles  only  are  equal.  Quite  a  number 
of  other  bodies  crystallize  in  the  same  forms.  All  the  known 
crystalline  forms  have  been  classified  into  six  systems,  and 
the  expert  crystallographer  is  able,  from  the  measurement  of 
its  angles  and  the  position  of  its  faces,  to  relegate  a  particular 
crystal  to  its  proper  system.  But  the  subject  is  a  wide  one, 
and  the  student  must  be  satisfied  at  this  stage  with  such 
information  as  he  can  get  by  an  examination  of  the  crystals 
he  produces. 

When  crystallization  takes  place  en  masse,  as  in  ordinary 
experiments,  the  individual  crystals  interfere  with  each  other, 


150  METALLUEGICAL  CHEMISTRY 

and  imperfect  forms  are  obtained  ;  but  such  parts  as  are  well 
defined  are  perfect  in  themselves,  and  serve  to  establish  the 
form.  To  obtain  a  perfect  crystal  of  considerable  size  it  is 
necessary  to  pick  out  a  small  perfect  one,  and  then  nurse  it  in 
a  saturated  solution  of  the  same  salt. 

Insoluble  Salts.— Some  salts  are  practically  insoluble  in 
water,  and  there  is  a  very  easy  way  of  forming  them  by 
taking  advantage  of  this  property.  A  solution  of  a  soluble 
salt  of  the  metal  is  simply  mixed  with  a  solution  of  the  acid, 
or  a  solution  of  a  soluble  salt  of  -the  acid  and  another  metal. 
The  possibility  of  the  formation  of  the  insoluble  salt  is  thus 
assured,  and  it  immediately  commences  to  separate  from  the 
solution  as  a  precipitate,  which  may  be  crystalline,  curdy,  or 
granular. 

EXP.  82. — Pour  a  few  cubic  centimetres  of  a  solution  of  silver 
nitrate  into  a  clean  test-tube  and  add  to  it  dilute  hydrochloric  acid, 
a  little  at  a  time,  with  vigorous  shaking  between  each  addition,  as 
long  as  a  precipitate  separates.  Repeat  the  experiment,  using  a 
solution  of  common  salt  in  place  of  the  hydrochloric  acid.  Filter, 
wash  and  dry  the  solid  salt. 

The  white  curdy  solid  is  silver  chloride,  and  its  formation 
is  expressed  by  the  equations  : 

AgN03  +  HC1  =  AgCl  +  HN03. 
AgN03  +  NaCl  =  AgCl  +  NaN03. 

Silver  chloride  is  not  soluble  in  nitric  acid. 

EXP.  83. — Repeat  the  last  experiment,  using  a  solution  of  barium 
chloride  and  dilute  sulphuric  acid,  or  a  solution  of  potassium  sulphate 
in  place  of  the  acid. 

The  white,  finely-divided  solid  is  barium  sulphate,  and  it  is 
insoluble  in  hydrochloric  or  nitric  acid  : 

BaCl2  +  H2S04  =  BaS04  +  2HC1. 
BaCl2  +  K2S04  =  BaS04  +  2KC1. 

These  reactions  are  used  as  tests  for  hydrochloric  and  sulphuric 
acids  and  their  soluble  salts. 


OXIDES,  ACIDS,  AND  SALTS  151 

Direct  Union  of  Elements  to  form  Salts.— Salts  of 
hydrochloric  acid  may  be  formed  by  the  direct  union  of  the 
metals  with  chlorine.  This  has  been  already  illustrated  in  the 
experiments  with  chlorine.  Thus,  when  chlorine  gas  is  passed 
over  finely-divided  silver,  the  chloride  of  the  metal  is  formed  : 

2Ag  +  Cl2  =  2AgCl. 

Bromine  and  iodine  act  similarly,  but  not  so  readily,  as 
chlorine. 

Direct  Union  of  Oxides  to  form  Salts. — In  the  case  of 
some  oxyacids  it  is  possible  to  form  their  salts  by  the  direct 
union  of  the  acid-forming  and  the  basic  oxides. 

Thus:  PbO  +  SOs  =  PbS04. 

Lead 
sulphate. 

=  CaC03. 

Calcium 
carbonate. 

=  Pb2Si04. 

Lead  silicate. 

These  reactions  indicate  that  there  cannot  be  any  strong 
objection  to  writing  the  formulae  of  the  compounds  in  such 
a  way  as  to  show  the  combination  of  the  two  oxides.  For 
example:  PbO.S03 ;  CaO.C02;  2PbO.Si02.  Sometimes  it 
is  convenient  to  use  such  formulae,  especially  in  the  case  of 
silicates.  See  Chap.  XIV. 

The  decomposition  of  a  carbonate  by  an  acid,  in  which  the 
salt  of  the  acid  and  of  the  metal  in  the  carbonate  is  formed, 
furnishes  another  general  method  of  forming  salts.  Illustra- 
tions of  this  method  will  be  found  in  Chap.  X.  The  following 
is  a  summary  of  the  chief  modes  of  salt  formation  :  (1)  Direct 
union  of  elements ;  (2)  direct  union  of  acid-forming  and  basic 
oxides ;  (3)  replacement  of  hydrogen  in  an  acid  by  a  metal ; 
(4)  reaction  between  acids  and  bases;  (5)  formation  of  in- 
soluble salts  by  precipitation  ;  (6)  decomposition  of  carbonates 
by  acids. 

Nomenclature  Of  Salts.— In  the  systematic  naming  of 


152  METALLUKGICAL  CHEMISTRY 

salts,  terminations  are  used,  as  in  the  case  of  oxides.  The 
assumption  that  acids  are  also  salts  of  the  displaceable  hydrogen 
they  contain,  and  are,  therefore,  of  the  same  general  character 
as  the  compounds  formed  by  the  replacement  of  such  hydrogen 
by  metals,  is  very  useful. 

According  to  this  view  hydrochloric  acid,  HC1,  is  also  hydro- 
gen chloride,  and  common  salt,  sodium  chloride.  Similarly 
we  have  bromides,  iodides,  etc.  Thus,  when  the  acid  and  its 
salts  contain  two  elements  only,  the  termination  -ide  is  used 
to  designate  the  group  to  which  the  salt  belongs,  and  the 
name  of  the  metal  to  particularize  the  salt.  Potassium 
bromide,  KBr;  zinc  iodide,  ZnI2;  bismuth  chloride,  BiCl3, 
are  examples.  Nitric  acid,  HN03,  is  also  hydrogen  nitrate, 
and  saltpetre  is  potassium  nitrate. 

But  there  is  also  another  acid,  HN02,  nitrous  acid  or 
hydrogen  nitrite,  and  its  salts  are  nitrites — e.g.,  potassium 
nitrite,  KNO2. 

Thus  when  there  are  two  acids  containing  the  same  elements, 
but  in  different  proportion,  the  terminations  -OUS  and  -ite  are 
used  for  the  lower  acid  and  its  salts,  and  -ic  and  -ate  for  the 
higher  compounds.  Similarly  sulphurous  acid  gives  sulphites, 
and  sulphuric  acid,  sulphates.  When  there  is  only  one  oxy- 
acid  of  a  given  element  the  higher  terminations  are  used  ; 
e.g.,  carbonic  acid  gives  carbonates.  The  student  is  advised 
to  keep  the  idea  of  acids  as  salts  of  hydrogen  well  in  mind,  as 
it  will  assist  him  in  getting  a  firm  grip  of  the  composition  and 
formulae  of  salts. 

Constitution  Of  Acids. — The  molecule  of  an  acid  may  be 
looked  upon  as  consisting  of  two  parts,  which  are  represented 
graphically  thus :  (H)(C1),  (H)(N08),  (H2)(S04).  Now,  the 
second  part  of  the  acid  molecule  is  present  in  the  molecule  of 
any  salt  of  the  acid,  and  is  characteristic  of  it ;  so  that  the 
(Cl),  (N03),  and  (S04)  groups  are  characteristic  of  chlorides, 
nitrates,  and  sulphates  generally.  The  same  may  be  said  of 
other  acids  and  their  salts. 


OXIDES,  ACIDS,  AND  SALTS  153 

There  is  good  reason  to  believe  that  when  an  acid  or  its 
salt  is  dissolved  in  a  very  large  quantity  of  water,  its  molecules 
split  up  or  dissociate  into  the  parts  mentioned  above.  These 
parts  are  called  ions  ;  and  even  when  the  solution  is  moderately 
strong,  an  average  number  of  molecules  are  believed  to  be 
constantly  in  this  state  of  dissociation,  or  ionic  condition. 
This  matter  is  closely  studied  in  electro-chemistry,  and  is  men- 
tioned here  in  order  to  justify  the  statements  made  in  the  last 
paragraph.  It  has  already  been  stated  that  some  metals  form 
two  basic  oxides,  and  when  these  dissolve  in  a  given  acid  two 
distinct  salts  are  formed,  Thus  there  are  two  sulphates  of 
iron  :  ferrous  sulphate,  FeS04,  and  ferric  sulphate,  Fe2(S04)3; 
also  two  chlorides  of  mercury,  mercurous  chloride,  Hg2Cl2, 
and  mercuric  chloride,  HgCl2.  But  it  must  not  be  inferred 
that  when  there  are  two  basic  oxides  of  a  metal  there  are  also 
two  salts  of  any  acid.  For  example,  the  two  oxides,  Cu20 
and  CuO,  both  give  copper  sulphate,  CuS04,  when  heated  with 
dilute  sulphuric  acid,  but  the  first  is  decomposed  and  leaves  a 
deposit  of  copper.  Thus  : 

Cu20  +  H2S04  =  CuS04  +  H20  +  Cu. 

There   are,    however,    two   chlorides   of   copper,    Cu2Cl2  and 
CuCl2. 

Some  old-fashioned  distinctive  names  which  are  still  used 
may  be  mentioned  here.  Thus  ferrous  sulphate  is  also 
the  protosulphate  of  iron,  and  ferric  sulphate  the  per- 
sulphate of  iron.  Also  subchloride  of  mercury  is  the  same 
as  mercurous  chloride,  and  the  perchloride  is  the  mercuric 
salt. 

Basicity  Of  Acids.— The  valency  of  metals  has  already 
been  defined  as  their  atom-replacing  power,  and  is  measured 
by  the  number  of  atoms  of  hydrogen  replaced  by  one  atom 
of  the  metal.  Now,  if  acids  are  salts  of  hydrogen,  it  is  easy 
to  see  from  their  formulae  that  the  number  of  atoms  of  the 
replaceable  hydrogen  differs  in  different  acids,  and  that  the 


154  METALLURGICAL  CHEMISTRY 

property  known  as  the  basicity  of  the  acid  depends  on  this 
number. 

Nitric  acid,  HN03    is  mono  basic. 

Sulphuric  acid,     H2S04   is  di  „    1 

Phosphoric  acid,  H3P04  is  tri  „     [-Poly  basic. 

Silicic  acid,  H4Si04  is  tetr        „    J 

With  a  monobasic  acid  the  hydrogen  is  all  replaced  in  one 
reaction,  and  the  normal  salt  only  is  formed.  But  with  a 
polybasic  acid  the  hydrogen  may  be  replaced  in  two  or  more 
distinct  reactions,  thus  giving  rise  to  two  or  more  distinct 
salts.  Those  in  which  hydrogen  is  present  are  called  acid 
salts. 

Potassium  hydrogen  sulphate,  KHS04,  acid  salt. 

Potassium  sulphate  K2S04,  normal  salt. 

Sodium  dihydrogen  phosphate,  NaH2P04,  acid  salt. 

Disodium  hydrogen  phosphate,  Na2HP04,  acid  salt. 

Trisodium  phosphate,  Na3P04,  normal  salt. 

The  second  salt  of  phosphoric  acid  above  mentioned  is  the 
common  sodium  phosphate.  Sometimes  salts  are  formed  in 
which  an  excess  of  the  acid-forming  oxide  is  present.  Such 
compounds  are  anhydro-add  salts.  The  well-known  example 
is  potassium  bichromate,  K2O207,  which  may  also  be  written 
K2Cr04Cr03.  Potassium  chromate,  K2Cr04,  is  the  normal 
salt. 

Similarly  some  salts  contain  an  abnormal  quantity  of  the 
basic  oxide  or  hydroxide,  and  are  then  called  basic  salts. 
The  green  mineral  malachite  is  a  basic  carbonate  of  copper, 
CuC03.CuH202. 

Only  one  acid  of  carbon  has  been  mentioned,  but  a  large 
number  of  acids  are  found  in  organized  matter,  both  animal 
and  vegetable,  or  are  formed  during  organic  changes.  These 
invariably  contain  carbon  and  hydrogen,  but  the  whole  of 
the  hydrogen  is  not  as  a  rule  replaceable  in  the  same  manner 
as  that  of  a  mineral  acid.  This  is  well  illustrated  in  the  case 
of  acetic  acid  and  its  salts.  The  empirical  formula  of  this 


OXIDES,  ACIDS,  AND  SALTS  155 

acid  is  H4C202,  but  to  show  its  basicity  it  must  be  written 
H.C2H302,  for  it  contains  only  one  atom  of  replaceable  hydro- 
gen which  is  written  first.  Thus  the  formula  for  lead  acetate 
(sugar  of  lead)  is  Pb(C2H302)2.  On  the  other  hand,  oxalic 
acid,  H2C204,  is  dibasic.  The  salts  of  these  acids  are  termed 
respectively  acetates  and  oxalates. 

Double  Salts. — It  is  found  that  when  some  soluble  salts 
are  dissolved  together  in  approximately  molecular  proportions 
they  also  crystallize  together  in  strictly  molecular  proportions 
when  the  solution  is  sufficiently  concentrated.  In  fact,  so 
constant  is  the  composition  of  these  double  compounds,  that 
they  are  regarded  as  definite  salts.  The  alums  are  the  most 
common  and  the  best  illustrations  of  these  very  interesting 
bodies.  Common  alum  is  a  double  sulphate  of  potassium  and 
aluminium,  and  its  empirical  formula  is  KA1(S04)2,12H2O. 
The  two  sulphates  are  potassium  sulphate,  K2S04,  and 
aluminic  sulphate,  A1_,(S04)3.  So  that  the  formula  which 
shows  the  composition  best  is  obtained  by  doubling  the  above 
thus,  K2Al2(S04)4,24HoO.  Potassium  sulphate  is  easily 
crj^stallized  alone,  but  the  aluminium  sulphate  only  with 
difficulty,  and  the  forms  of  the  crystals  are  quite  distinct.  On 
dissolving  the  salts  together  in  molecular  proportions,  and 
allowing  the  solution  to  crystallize,  the  double  salt  is  obtained 
in  crystals  of  the  regular  octahedron  form. 

EXP.  84.—  Put  1  gram  of  alumina,  Al.203j  into  a  small  beaker,  add 
10  c.c.  of  dilute  sulphuric  acid  (1  to  3),  and  boil  slowly  until  the 
solid  disappears,  and  a  somewhat  milky-looking  solution  is  obtained. 
Dissolve  1  gram  of  caustic  potash  in  10  c.c.  of  water,  add  this  solu- 
tion to  the  first,  raise  the  whole  to  boiling,  and  filter  the  solution 
into  a  porcelain  basin.  Set  it  aside  to  crystallize.  Examine  the 
crystals  under  a  lens,  and  note  their  characteristic  shape.  Set  aside 
the  best  for  reference. 

The  crystals  are  usually  well  defined,  if  the  solution  is  not 
too  concentrated,  and  the  crystallization  takes  place  slowly. 
If  caustic  soda  is  used  in  the  above  experiment  in  place  of 
potash,  exactly  similar  crystals  are  obtained.  The  two  salts 


156  METALLURGICAL  CHEMISTRY 

cannot  be  distinguished  from  each  other  by  their  crystalline 
form.  Also,  if  ferric  oxide,  Fe203,  is  substituted  for  alumina, 
light  green  crystals  are  obtained,  which  cannot  be  distinguished 
from  the  others  by  their  crystalline  form. 

EXP.  85. — Dissolve  5  grams  of  potassium  bichromate  in  25  c.c.  of 
water,  and  add  1  c.c.  of  strong  sulphuric  acid.  Prepare  sulphur 
dioxide,  and  pass  the  washed  gas  (p.  95)  through  the  solution  until 
it  turns  deep  green  in  colour.  Pour  it  into  a  large  porcelain  dish, 
and  set  it  aside  to  crystallize.  Compare  the  crystals  of  chrome  alum 
with  those  of  common  alum  from  Exp.  84. 

It  will  be  noticed  that  the  two  metals  in  the  sulphates  are 
monad  and  triad  respectively ;  any  similar  metal  may  be 
substituted  in  the  compound  without  altering  its  form.  Thus 
silver  and  the  monad  group  ammonium  (NH4)  may  take  the 
place  of  potassium,  and  the  triad  metals,  chromium  and 
manganese,  that  of  aluminium.  The  general  formula  is 
M'M'"(S04)2,12H20  where  M'  and  M'"  are  monad  and  triad 
atoms  respectively. 

If  a  crystal  of  common  alum  is  put  into  a  solution  of  iron 
alum  it  increases  in  bulk ;  the  enlarged  crystal  may  then  be 
transferred  in  turn  to  solutions  of  chrome  alum  and  manganese 
alum,  in  each  of  which  there  is  a  further  growth,  but  the  final 
crystal  keeps  its  fundamental  form.  Further,  if  a  mixture  of 
various  alums  is  crystallized,  one  crop  of  crystals  containing 
the  different  compounds  is  obtained. 

Another  class  of  double  salts  of  practical  importance  contain 
ammonium  or  potassium  sulphate  with  iron,  nickel,  or  zinc 
sulphate.  The  well-known  examples  are  ferrous  ammonium 
sulphate  (NH4)2Fe(S04)2,6H20,  and  nickel  ammonium  sul- 
phate, (NH4)2NiS04,6H20. 

Isomorphism. — The  similarity  of  form  of  different  crystal- 
lized compounds  is  looked  upon  as  an  indication  of  similarity 
in  the  constitution  of  their  molecules,  and  such  compounds  are 
said  to  be  isomorphous,  which  means,  literally,  of  "  equal 
form."  It  is  not  confined  to  the  double  salts  described  above, 
for  two  single  salts  crystallized  separately  may  have  the  same 


OXIDES,  ACIDS,  AND  SALTS  157 

crystalline  form.     They  are  then   judged  to  be  made  up  of 
similarly  constituted  molecules. 

This  supposed  similarity  of  the  molecules  of  isomorphous 
compounds  suggests  similarity  between  the  atoms  of  those 
elements  which  are  able  to  replace  each  other  in  them  without 
altering  the  crystalline  form  of  the  compounds ;  and  they  are 
also  considered  to  be  isomorphous.  Thus,  Na,  K,  Ag ;  and  Al'", 
Fe'",  Cr'"  replace  each  other  in  the  alums.  They  are  regarded 
as  two  groups  of  isomorphous  metals.  Similarly,  Zn",  Fe",  Ni" 
are  isomorphous  with  regard  to  the  class  of  isomorphous  salts 
of  which  (NH4)2Ni(S04)2,6H20  is  a  well-known  example. 
This  principle  of  isomorphism  has  been  of  considerable  value 
in  helping  to  fix  the  atomic  weights  of  a  number  of  the 
metals.  It  has  also  been  used  to  explain  some  of  the  changes 
which  take  place  during  the  cooling  and  solidification  of  a 
molten  mass  of  mixed  metals.  A  mixture  of  isomorphous 
salts  cannot  be  separated  from  each  other  by  crystallization, 
as  they  are  deposited  from  the  solution  together ;  but  mixtures 
of  non-isomorphous  compounds  may  be  so  separated  if  their 
solubility  in  the  mother  liquor  differs.  Similar  remarks 
apply  to  isomorphous  and  non-isomorphous  metals  in  a  molten 
alloy. 

Action  Of  Heat  upon  Salts.— Water  of  crystallization  is 
driven  off,  usually  at  temperatures  very  little  above  the  boil- 
ing-point, 100°  C.  But  sometimes  one  or  more  molecules  of 
water  hold  more  tenaciously  to  the  salts,  and  require  a  higher 
temperature  for  their  separation.  Further  heating  of  the 
anhydrous  salt  may  cause  it  to  melt  and  even  to  volatilize ;  or 
it  may  be  decomposed  into  simpler  compounds,  and,  in  a  few 
cases,  into  its  elements. 

EXP.  86. — Weigh  exactly  2  grams  of  powdered  crystallized  copper 
sulphate  in  a  dry  porcelain  crucible,  or  on  a  clay  roasting  dish,  and 
heat  it  on  a  pipeclay  triangle  over  a  Bunsen  flame  until  it  turns 
white.  Allow  the  crucible  to  cool  and  reweigh  it.  Put  the  crucible 
into  a  gas  muffle,  and  keep  it  at  a  fair  red  heat  for  twenty  minutes, 


158  METALLURGICAL  CHEMISTRY 

or  until  the  residue  is  quite  black.  Allow  the  crucible  to  cool  and 
reweigh  it.  The  residue  is  black  oxide  of  copper. 

EXAMPLE.—  Weight  of  crucible  +  salt    =13140  <  _n  71ft 
„      after  first  heating  =  12  -422  j  ~ 

„    second  „       =  11-776     =  0'646 
„     of  crucible  =11-140     =  0'636 

First  loss  =  0-718  gram  H20;  second  loss  =  0'646  gram  S03  ; 
residue  =  0'636  gram  CuO,  and  these  are  the  weights  of  the 
three  oxides  in  2  grams  of  the  compound.  Now,  divide  each 
weight  by  the  molecular  weight  of  the  oxide  to  which  it  belongs. 
CuO  =  79;  S03  =  80;  H20  =  18. 


Then,  -  0-0399;  0-008;  =  0-008. 

18  oU  79 

These  ratios  show  the  relation  between  the  molecules  of  the 
oxides  in  the  crystallized  salt,  and  are  readily  reduced  to 
whole  numbers  by  dividing  each  by  the  lowest.  This  gives 
1  CuO  :  1  S03  :  5  H^O,  and  the  formula  may  be  written, 
CuO.S03.5H20.  The  complete  reaction  is  expressed  by  the 
equation  : 

CuS04,5H20  =  CuO  +  S08  +  5H20. 

Four  molecules  of  water  are  separated  at  a  temperature  about 
100°  C.  ;  but  the  last  molecule  clings  until  the  temperature 
exceeds  200°  C.  A  similar  experiment  may  be  made  with 
crystallized  ferrous  sulphate,  FeS04,7H20,  which  gives  the 
following  reaction  when  heated  : 

2(FeS04,7H20)  =  Fe203  +  S02  +  S08  +  14H2O. 

Most  of  the  sulphates  of  the  common  heavy  metals  are  decom- 
posed in  a  similar  manner,  but  a  number  of  others,  including 
those  of  potassium,  barium,  lead,  etc.,  may  be  heated  very 
strongly  without  undergoing  decomposition. 

The  Nitrates  are  all  more  or  less  readily  decomposed  when 
heated.  Oxides  of  nitrogen  and  oxygen  are  set  free,  and  a 
residue  of  the  oxide  of  the  metal  left.  Nitrates  of  silver  and 
mercury  are  reduced  to  the  metallic  state. 


OXIDES,  ACIDS,  AND  SALTS 


159 


The  Chlorides  fuse  more  or  less  readily,  and  volatilize  when 
strongly  heated.  Some  are  decomposed  with  liberation  of 
chlorine,  and  are  thus  reduced,  either  to  a  lower  chloride  or  to 
the  metallic  state. 

The  Carbonates  of  sodium  and  potassium  will  withstand 
a  high  temperature  without  decomposition.  The  other  common 
carbonates  are  split  up  into  metallic  oxides  and  carbon  dioxide. 
Barium  carbonate  requires  a  bright  red  heat  for  its  decom- 
position. 

A  number  of  well-known  acids  are  represented  by  only  a 
few  common  salts,  but  as  these  are  of  considerable  importance 
a  list  of  them  is  given  below. 


ACID. 

FORMULA. 

SALT. 

FORMULA. 

COMMON  NAME. 

Sulphurous 
Thio- 

H2S03 
H2SA 

Sodium  sulphite 
Sodium  thio- 

Na2S03,10H20   'Sulphite  of  soda 
Na2S,03.5H20      Hyposulphite 

sulphuric 

sulphate 

A       4         O                   &                                              «                1 

of  soda 

Phosphoric 

H3P04 

Disodium  hydro- 

Na2HP04,12H20    Phosphate  of 

gen  phosphate 

soda 

Chloric 

HC103 

Potassium 

KC103               Chlorate  of 

chlorate 

potash 

Chromic 

H2Cr04 

Potassium 

K2Cr04             Chromate  of 

chromate 

potash 

Potassium 

K2Cr207 

bichromate 

Permanganic 

H2Mn208 

Potassium  per- 

K2Mn208 

manganate 

Metaboric 

HB02 

Acid  sodium 

Na2B407,10H20  \         Borax 

borate 

Nitrous 

HN02 

Potassium  nitrite 

KN02 

Nitrite  of 

potash 

Hydro- 

HBr 

Potassium 

KBr 

bromic 

bromide 

Hydriodic 

HI 

Potassium  iodide 

KI 

Hydrofluoric 

HF 

Calcium  fluoride 

CaF2 

Fluor  spar 

Salts  of  these  acids  and  the  common  metals  are  easily 
obtained  by  one  or  other  of  the  usual  methods  of  salt 
formation. 

SUMMARY. 

The  basic  oxides  of  the  metals  are  a  very  important  class  of 
compounds.  They  react  with  acids  to  form  salts.  In  the 


160 


METALLUEGICAL  CHEMISTEY 


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162  METALLURGICAL  CHEMISTRY 

majority  of  cases  the  basic  character  of  the  oxide  disappears 
along  with  the  acid  character  of  the  acid,  and  neutral  bodies 
are  formed.  Metallic  hydroxides  contain  the  elements  of 
water  in  addition  to  the  oxides.  They  act  in  the  same  way 
towards  acids  as  the  corresponding  oxides.  The  peroxides 
undergo  partial  decomposition  when  heated,  and  oxygen  is 
liberated.  The  acid-forming  oxides  unite  with  water  to  form 
acids,  and  in  this  respect  they  resemble  the  acid-forming 
oxides  of  non-metals.  Salts  are  formed  in  a  variety  of  ways, 
but  a  given  salt  always  has  the  same  composition,  however 
it  may  be  formed.  The  quantitative  relations  of  neutraliza- 
tion are  just  as  definite  as  those  of  any  other  mode  of  chemical 
action. 

QUESTIONS. 

1.  How  are  oxides  usually  classified1?     Give  examples. 

2.  What  are  the  general  effects  of  adding  basic  oxides  to 
acid  solutions  1 

3.  Explain  how  an  acid  salt  and  a  normal  salt  differ  from 
each  other. 

4.  How  are  crystals  usually  formed  1     Why  are  some  salts 
said  to  be  isomorphous  1 

5.  What  is  meant  by  the  nomenclature  of  compounds  1   Give 
examples  to  illustrate  your  answer. 

6.  Describe  three  ways  of  forming  zinc  sulphate. 

7.  Is  there  any  similarity  between  the  basicity  of  an  acid  and 
the  valency  of  a  metal  ? 

8.  Give  a  short  account  of  the  principle  of  isomorphism. 


CHAPTER  X 
CAEBON  AND  ITS  COMPOUNDS 

THE  element  carbon  is  a  very  common  and  useful  substance. 
It  is  a  constituent  of  all  animal  and  vegetable  bodies;  and 
is  present  among  the  rocks  and  minerals  of  the  earth's  crust 
in  all  natural  carbonates,  of  which  the  calcium  salt  CaC03 
is  by  far  the  most  abundant.  It  is  also  the  principal  con- 
stituent of  coal,  which  is  the  product  of  the  very  slow  decay 
of  vegetable  matter  out  of  contact  with  air.  Such  matter 
contains  carbon,  hydrogen,  oxygen,  and  nitrogen  as  essential 
constituents,  together  with  a  small  quantity  of  inorganic 
or  earthy  matter.  Air-dried  wood  contains  about  50  per  cent, 
of  carbon  and  1*5  per  cent,  of  inorganic  compounds;  the 
remainder  is  made  up  of  oxygen,  hydrogen,  and  nitrogen. 
During  the  extremely  slow  changes  which  the  original  wood 
has  undergone  in  its  conversion  to  coal,  the  oxygen  and 
hydrogen,  together  with  some  of  the  carbon,  have  been  largely 
eliminated.  Therefore  the  residue  of  coal  contains  a  much 
higher  percentage  of  carbon  than  the  vegetation  from  which 
it  was  formed.  In  fact,  the  oldest  variety  of  coal  may  be 
described  as  impure  carbon,  containing  as  it  does  upwards  of 
90  per  cent,  of  carbon. 

Graphite  or  plumbago  is  also  an  impure  form  of  carbon  found 
in  the  earth's  crust ;  and  the  diamond  is  the  purest  form  of 
natural  carbon. 

The  various  oils  and  fats  of  animal,  vegetable,  and  mineral 
origin  are  extremely  rich  in  carbon. 

163  11—2 


164  METALLURGICAL  CHEMISTRY 

Charcoal- — When  wood  is  heated  in  a  vessel  from  which 
the  air  is  excluded,  such  as  a  porcelain  crucible  with  the  lid 
on,  a  residue  very  rich  in  carbon  is  obtained;  it  is  called 
charcoal,  and  is  a  very  useful  fuel. 

EXP.  87. — Heat  a  piece  of  sound  wood,  about  ^-inch  cube,  in  a 
porcelain  crucible  with  the  lid  on  until  gaseous  matter  ceases  to 
escape.  The  crucible  should  be  put  on  a  pipeclay  triangle  over  a 
Bunsen  flarne  for  a  few  minutes,  and  the  heating  finished  over  the 
foot  blow-pipe.  When  the  crucible  is  nearly  cold,  remove  the  lid 
and  examine  the  residue. 

Grind  the  piece  of  charcoal  to  a  fine  powder  in  a  clean  mortar, 
put  it  into  a  small  platinum  crucible  with  a  closely-fitting  lid,  and 
heat  it  as  strongly  as  possible  for  ten  minutes  in  the  blow-pipe 
flame. 

The  experiment  may  be  made  quantitative  by  weighing  the  wood 
before,  and  the  charcoal  after,  heating. 

The  charcoal  thus  obtained  contains  about  92  per  cent,  of 
carbon,  1*5  per  cent,  of  earthy  matter,  which  forms  the  ash 
when  the  charcoal  is  completely  burnt,  and  some  oxygen  and 
hydrogen. 

Some  carbonaceous  bodies  are  practically  free  from  earthy 
matter,  and  furnish  very  pure  charcoal  when  properly  treated. 
Sugar  is  a  very  good  illustration  of  these  bodies. 

EXP.  88.— Melt  some  good  lump-sugar  in  a  large  porcelain 
crucible ;  continue  the  heating  until  the  combustible  gas  is  all 
expelled,  and  a  black  residue  of  "  burnt  "  sugar  is  obtained.  The 
total  quantity  of  sugar  should  be  added  to  the  crucible  a  little  at  a 
time  as  it  works  down,  or  it  will  boil  over  and  make  the  experiment 
unsatisfactory. 

Grind  the  black  residue  to  a  fine  powder,  and  heat  it  very  strongly 
for  ten  minutes  in  a  closed  platinum  crucible,  as  described  above. 
When  cold,  transfer  the  powder  to  a  small  stoppered  bottle. 

The  carbon  so  obtained  is  not  quite  pure,  but  contains 
about  0*1  per  cent,  of  earthy  matter  and  some  oxygen  and 
hydrogen.  The  gaseous  elements  can  be  got  rid  of  by  strongly 
heating  the  powder  in  a  glass  tube  through  which  a  stream  of 
dry  chlorine  gas  is  made  to  pass.  The  heating  in  contact  with 
chlorine  is  continued  as  long  as  fumes  of  hydrochloric  acid 
escape  from  the  open  end  of  the  tube.  A  sample  of  carbon 


CARBON  AND  ITS  COMPOUNDS  165 

prepared  as  described  in  Exp.  88,  and  also  subjected  to  the 
action  of  chlorine,  was  found  to  contain  99*7  per  cent,  of 
carbon  and  O'l  per  cent,  of  ash. 

COMBUSTION  OF  CHARCOAL. 

When  charcoal  is  heated  in  a  good  current  of  air  or  in 
oxygen,  it  burns,  and  the  whole  of  the  carbon  in  it  is  converted 
into  carbon  dioxide,  C02.  This  compound  is  a  colourless, 
transparent  gas,  which  dissolves  in  water,  and  forms  a  weak 
acid  solution.  It  is  therefore  an  acid-forming  oxide,  and 
combines  with  a  number  of  basic  oxides  to  form  the  salts 
known  as  carbonates.  It  is  rapidly  absorbed  by  a  solution 
of  caustic  soda  or  potash,  and  forms  a  carbonate  with  the 
alkali  which  remains  in  the  solution.  The  reaction  is  shown 
by  the  equation : 

2KHO  +  C02  =  K2C03  +  H20. 

Potassium 
carbonate. 

This  change  is  taken  advantage  of  in  estimating  how  much 
carbon  dioxide  is  formed  by  the  combustion  of  a  known  weight 
of  carbon.  The  experiment  requires  considerable  care,  and 
accurate  weighings  must  be  made ;  but  the  following  descrip- 
tion will  be  readily  followed. 

EXP.  89. — A  small  quantity  (about  O'l  gram)  of  the  purest  char- 
coal at  hand  is  weighed  on  a  bent  strip  of  platinum  foil,  which  is 
then  pushed  into  the  middle  of  a  glass  combustion-tube,  A  (Fig.  30). 
One  end  of  A  is  then  connected  with  the  weighed  potash  bulbs,  B. 
which  contain  a  strong  solution  of  caustic  potash,  KHO,  in  water 
(1  to  2).  The  other  end  of  A  is  connected  with  the  drying-tube  C, 
containing  small  lumps  of  caustic  soda  in  the  limb  a,  and  calcium 
chloride  in  the  limb  b.  The  part  of  the  tube  directly  under  the 
platinum  foil  containing  the  charcoal  is  then  heated  in  a  Bunsen 
flame,  and  a  moderate  stream  of  dry  oxygen  passed  through  the 
apparatus.  This  is  readily  effected  by  using  the  apparatus  shown  in 
Fig.  32.  The  bottle  C  is  filled  with  oxygen,  and  the  wash-bottle  D, 
containing  a  little  strong  sulphuric  acid,  is  connected  with  the  inlet 
of  the  drying-tube  C  (Fig.  30).  The  charcoal  burns  in  the  oxygen  ; 
the  carbon  dioxide  formed  is  carried  forward  by  the  excess  of  gas, 
and  is  completely  absorbed  by  the  potash  in  B.  The  difference  in 


166 


METALLUKGICAL  CHEMISTRY 


the  weight  of  the  platinum  foil  containing  the  charcoal,  before  and 
after  the  combustion,  gives  the  weight  of  carbon  burnt,  and  the 
difference  in  the  weight  of  the  potash  bulbs  gives  the  weight  of 
carbon  dioxide  absorbed.  A  small  quantity  of  ash  is  left  on  the 
foil. 


FIG.  30. 


For  very  accurate  work  a  drying-tube  must  be  placed 
between  A  and  B.  Powdered  graphite  and  diamond  dust 
burai  very  slowly  in  a  glass  tube ;  but  a  porcelain  tube  heated 
in  a  table  furnace  (Fig.  8)  gives  very  satisfactory  results. 
When  pure  carbon  is  used  the  ratio,  carbon  to  carbon 
dioxide,  is  found  to  be  12  :  44.  That  is,  12  parts  of  carbon 
combine  with  32  parts  of  oxygen  to  form  44  parts  of  carbon 
dioxide.  Therefore  its  formula  is  C02' 

Experiments  made  with  sugar  carbon,  purified  graphite, 
and  diamond  dust  give  practically  the  same  result,  and  either 
of  them  can  be  used  to  determine  the  composition  of  carbon 
dioxide. 

Experimental  proof  is  thus  furnished  of  the  statement  that 
charcoal,  graphite,  and  diamond  are  one  and  the  same 
Chemical  element — i.e.,  carbon.  But  there  is  a  very  marked 
difference  in  their  physical  properties.  These  variations  in 
the  properties  of  the  same  element  are  called  allotropic 
modifications  (p.  57).  The  principal  variations  are  in  the 
colour,  hardness,  and  structure  of  the  bodies.  The  diamond 
is  crystalline  and  very  hard ;  graphite  is  crystalline  and  soft ; 


CABBON  AND  ITS  COMPOUNDS  167 

charcoal  is  amorphous — i.e.,  without  crystalline  form.  Another 
important  difference  is  that,  although  the  equal  weights  of  the 
three  give  the  same  weight  of  carbon  dioxide,  they  do  not 
give  out  the  same  quantity  of  heat  when  burnt.  The  calorific 
power  of  the  diamond  is  less  than  that  of  graphite,  and  of 
graphite  less  than  that  of  charcoal.  This  is  probably  due  to 
the  constitution  and  aggregation  of  their  molecules.  The 
different  forms  of  carbon  can  be  changed  from  one  to  the 
other.  1  hus,  when  a  diamond  is  raised  to  the  temperature  of 
the  electric  arc,  it  is  converted  into  a  black  mass  of  coke -like 
carbon.  Amorphous  carbon,  after  absorption  by  heated  iron, 
may  be  caused  to  separate  from  the  molten  mass  in  the  form 
of  graphite.  This  form  of  artificial  graphite  is  known  as  kisk. 
Scales  of  crystalline  graphite  are  easily  recognised  on  the 
fractured  surface  of  a  piece  of  grey  pig-iron.  The  carbon 
goes  into  the  furnace  in  the  form  of  coke.  Amorphous  carbon 
has  also  been  converted  into  diamond  crystals ;  but  those 
produced  up  to  the  present  are  too  small  to  be  of  practical  use. 

CARBONATES. 

Native  Carbonates.  — Limestone,  marble,  and  chalk, 
which  occur  in  enormous  quantities  in  the  earth's  crust,  are 
varieties  of  calcium  carbonate,  CaC03.  The  first  two  are  dis- 
tinctly crystalline  in  character.  Calcite,  or  Iceland  spar,  a 
purer  form  of  the  same  compound,  is  found  in  the  form  of 
rhombic  crystals,  which  are  sometimes  as  clear  and  trans- 
parent as  glass.  Aragonite  is  another  crystalline  form  of 
calcium  carbonate. 

Magnesium  carbonate,  MgC03,  if  isomorphous  with  the 
calcium  salt,  and  occurs  along  with  it  in  dolomite,  which  is 
usually  described  as  a  double  carbonate  of  lime  and  magnesia. 
Different  specimens  of  dolomite  contain  the  two  carbonates  in 
varying  proportions,  which  is  consistent  with  the  formation  of 
isomorphous  mixtures  (Chap.  IX.). 

Ferrous  carbonate,  FeC03,  which  is  the  principal  constituent. 


168  METALLUEGICAL  CHEMISTEY 

in  spathic  iron  ore  and  clay  ironstone,  and  manganous  car- 
bonate, MnC03,  are  isomorphous  with  the  calcium  and 
magnesium  salts  which  often  accompany  them  in  their  ores. 

Zinc  carbonate,  ZnC03,  in  calamine,  and  lead  carbonate, 
PbC03,  in  cerussite,  are  also  important  natural  carbonates. 

Prepared  Carbonates.— Sodium  carbonate,  Na2C03,  is 
the  most  important  of  these  compounds.  It  is  prepared  both 
in  the  anhydrous  form,  Na2CO3,  and  with  water  in  soda 
crystals,  Na2C03,10H20.  Both  forms  are  readily  soluble  in 
water.  Soda  crystals  lose  water  when  exposed  to  the  air, 
and  gradually  break  down  to  a  white  powder ;  they  are 
efflorescent.  When  heated  the  crystals  melt  in  their  water 
of  crystallization,  which  is  then  removed  by  evaporation, 
and  the  anhydrous  salt  is  left  as  a  white  residue.  This  melts 
to  a  clear  liquid  at  a  moderately  high  temperature.  If  mixed 
with  the  corresponding  potassium  salt  a  more  readily  fusible 
mass  is  obtained.  Common  fusion  mixture  contains  Na2C03  + 
K2C03. 

The  majority  of  metallic  carbonates  are  either  insoluble  or 
only  slightly  soluble  in  water,  and  are  precipitated  as  amorphous 
solids  when  a  solution  of  sodium  or  potassium  carbonate  is 
added  to  solutions  of  the  soluble  compounds  of  the  metals. 
Very  often,  however,  the  oxide  of  the  metal  separates  at  the 
same  time,  and  basic  carbonates  are  formed. 

EXP.  90. — Make  a  solution  of  sodium  carbonate  (1  in  10).  Pour 
about  10  c.c.  of  a  solution  of  calcium  chloride,  CaCl2.  into  a  large 
test-tube,  and  add  the  sodium  carbonate  solution  a  little  at  a  time, 
with  shaking,  as  long  as  a  precipitate  forms.  Kaise  the  contents  of 
the  tube  to  boiling,  filter,  and  well  wash  the  precipitate  on  the  filter. 
Dry  the  filter,  and  scrape  off  the  white  solid  calcium  carbonate. 
Repeat  the  experiment  with  a  solution  of  zinc  sulphate  in  place  of 
the  calcium  chloride.  Put  some  of  the  calcium  carbonate  in  a  test- 
tube,  add  a  little  dilute  hydrochloric  acid,  and  pour  the  escaping  gas 
into  a  test-tube  containing  a  little  lime-water.  Shake  the  lime-water 
to  bring  it  in  contact  with  the  gas,  and  note  the  milkiness  caused  by 
the  absorption  of  carbon  dioxide. 

Put  a  little  of  the  dry  zinc  carbonate  into  a  small  test-tube, 
heat  it  in  the  Bunsen  flame,  and  prove  that  carbon  dioxide  is 
evolved. 


CAEBON  AND  ITS  COMPOUNDS  169 

The  reactions  for  the  formation  of  the  carbonates  are  as 
follows  : 

CaCl9  +  Na2C03  =  CaC03  +  2NaCl. 
ZnSO*  +  Na2C03  =  ZnC03  +  Na2S04. 

Bicarbonates. — These  are  acid  salts,  of  which  the  best 
example  is  bicarbonate  of  soda,  NaHC03.  The  normal  soluble 
carbonates  have  a  decidedly  alkaline  reaction,  for  the  weak, 
acid-forming  oxide,  C09,  does  not  neutralize  the  strong  soda 
base,  Na20.  The  acid  carbonates  are  milder  because  of  the 
excess  of  C02  present.  They  are  readily  decomposed  at  a 
moderate  heat  into  the  normal  salts,  carbon  dioxide,  and 
water.  Normal  carbonates  are  converted  into  the  correspond- 
ing acid  salts  when  exposed  to  the  action  of  carbon  dioxide 
and  water,  and  sometimes  an  insoluble  carbonate  is  thus  con- 
certed into  a  soluble  one.  When  carbon  dioxide  is  passed  for 
some  time  into  a  solution  of  sodium  carbonate  the  bicarbonate 
crystallizes  out : 

Na2C03  +  C02  +  H20  =  2NaHC03. 

And  when  the  gas  is  passed  through  water  in  which  finely- 
divided  calcium  carbonate  is  suspended,  the  solid  passes  into 
solution.  The  acid  carbonate  is  soluble  : 

CaC03  +  C02  +  H20  =  CaH2(C03)2. 

Acid  carbonate." 

But  on  boiling  the  solution  the  C02  is  expelled  and  the  normal 
carbonate  precipitated. 

ACTION  or  HEAT  ON  CARBONATES. 

The  sodium  and  potassium  compounds  are  not  decomposed 
at  a  bright-red  heat.  Carbonates  of  Ba,  Sr,  Ca,  and  Mg  are 
decomposed  at  a  good  red  heat  into  the  basic  oxides  of  the 
metals  and  carbon  dioxide;  the  barium  carbonate  requires 
the  highest  temperature.  Carbonates  of  the  other  common 
metals  are  decomposed  at  a  still  lower  temperature.  Bicar- 
bonates give  off  water  in  addition  to  carbon  dioxide,  but  in 


170  METALLUBGICAL  CHEMISTKY 

the  case  of  the  alkaline  bicarbonates  only  part  of  the  CO2  is 
liberated  : 

2NaHC03  =  Na2C03  +  C02  +  H20. 

EXP.  91.  —  Weigh  1  gram  of  ground  marble  in  a  weighed  porcelain 
crucible  ;  heat  the  crucible,  without  the  lid,  in  a  gas  muffle  to  a 
bright-red  heat  for  fifteen  minutes.  When  the  crucible  is  cold  re- 
weigh  it.  The  loss  in  weight  gives  the  carbon  dioxide  driven  off. 
If  a  muffle  cannot  be  used,  heat  the  crucible  over  a  good  circular 
burner,  and  finish  with  five  minutes  over  the  foot  blow-pipe.  In  this 
case  a  platinum  crucible  is  best. 

EXAMPLE.—  One  gram  of  marble  lost  0  434  gram  on  being  heated 
for  twenty  minutes  in  a  gas  muffle.  Compare  this  with  the 
equation  : 

CaC0  = 


Sodium  bicarbonate  is  readily  decomposed  in  a  porcelain 
crucible  heated  over  a  Bunsen  burner,  and  as  the  commercial 
salt  is  comparatively  pure,  a  good  result  is  obtained.  Zinc 
and  lead  carbonates  also  furnish  good  examples.  The  residue 
of  zinc  oxide  is  quite  infusible  ;  but  lead  oxide  is  somewhat 
fusible,  and,  unless  the  temperature  is  carefully  regulated,  it 
will  soften,  stick  to  the  crucible,  and  spoil  it. 

When  a  basic  carbonate  containing  an  excess  of  the  an- 
hydrous oxide  is  heated  C02  only  is  driven  off,  but  when  the 
hydroxide  forms  the  excess  of  the  base  H20  is  also  evolved. 
In  this  case  the  loss  includes  both  carbon  dioxide  and  water. 
In  order  to  determine  the  proportions  of  the  three  oxides  in 
the  salt  it  is  necessary  to  determine  the  carbon  dioxide  by  a 
separate  experiment,  and  then  deduct  its  weight  from  the  total 
loss  on  heating. 

EXP.  92.  —  Weigh  1  gram  of  blue  copper  carbonate  in  a  weighed 
porcelain  crucible,  and  heat  it  over  the  Bunsen  flame  until  it  turns 
completely  black.  Allow  the  crucible  to  cool,  and  reweigh  it.  The 
loss  gives  the  weight  of  carbon  dioxide  and  water  driven  off. 

EXAMPLE.  —  One  gram  of  the  blue  carbonate  lost  0-296  gram  when. 
heated  for  ten  minutes  over  the  Bunsen  flame. 


CAEBON  AND  ITS  COMPOUNDS  171 

CARBONATES  AND  ACIDS. 

All  carbonates  are  decomposed  by  dilute  acids  with  evolution 
of  carbon  dioxide,  and  formation  of  the  salts  of  the  metals  in 
the  carbonates  and  of  the  acids.  If  the  salt  formed  is  insoluble, 
or  only  slightly  soluble,  in  the  acid  solution,  the  action  soon 
ceases  as  the  carbonate  becomes  coated  with  this  insoluble 
body,  which  protects  it  from  the  acid.  Nitric  acid  is  the  best 
general  solvent. 

Determination  of  the  Weight  of  C02  in  a  Car- 
bonate.— The  apparatus  shown  in  Fig.  27  gives  excellent 
results,  and  is  easy  of  manipulation.  It  may  be  used  for  any 
carbonate. 

EXP.  93.— Add  10  c.c.  of  dilute  nitric  acid  (1  to  8)  to  the  flask  A 
(Fig.  27),  and  fit  the  parts  of  the  apparatus  together.  Put  the  whole 
on  the  scale-pan  of  the  balance,  and  along  with  it  a  piece  of  filter- 
paper  which  will  conveniently  hold  1  gram  of  the  salt.  Weigh  the 
whole  accurately,  then  put  a  gram  weight  on  the  weight-pan,  and 
weigh  1  gram  of  blue  copper  carbonate  on  the  filter-paper.  Fold 
up  the  paper  containing  the  salt  into  a  packet  which  can  be  readily 
pushed  through  the  neck  of  the  flask.  Take  the  bung  out ;  remove 
the  stopper  from  the  outlet  of  the  drying-tube  ;  push  the  packet 
into  the  neck  of  the  flask,  and  then  into  the  body  with  the  tube 
passing  through  the  bung,  rapidly  inserting  the  latter  as  the  packet 
falls  in  the  acid.  When  the  action  is  finished,  connect  the  outlet  of 
the  drying-tube  with  the  aspirator  (Fig.  28),  and  draw  about  a  litre 
of  dry  air  through  the  flask  to  displace  the  remaining  C02  still 
present,  and  to  restore  the  atmosphere  in  the  flask  to  its  original 
condition.  Eeplace  the  stoppers  and  reweigh  the  apparatus.  The 
loss  gives  the  weight  of  C02  liberated  from  1  gram  of  the  salt,  for 
nothing  else  escapes  from  the  apparatus. 

EXAMPLE  :  Weight  of 

apparatus  +  filter-paper +  1  gram  of  the  salt  =  77v54  grams 
,,  ,,  ,,        after  experiment  =  77*449      ,, 

Loss  due  to  escape  of  dry  C02  =    0'205   gram 

From  the  data  supplied  by  Experiments  92  and  93  it  is 
found  that  1  gram  of  blue  copper  carbonate  contains  O704 
gram  of  CuO,  O205  gram  of  C02,  and  O091  gram  of  H20. 
Then  by  using  the  molecular  weights,  CuO -=79,  C02  =  44, 


172  METALLUKGICAL  CHEMISTKY 

H20  =  18,  the  ratio  between  the  oxides  in  the  carbonate  can 
be  found  and  the  formula  deduced.  Thus  : 

O^l=0.0891;  0^,0-00465;  ^=0.005. 

By  reducing  these  quotients  to  the  nearest  whole  numbers 
the  ratio  2  :  1  :  1  is  obtained,  and  the  empirical  formula  of 
the  compound  is  2CuO.C02.H20,  which  is  thus  proved  to  be 
a  basic  carbonate.  The  composition  of  the  basic  carbonates  of 
lead  and  zinc  may  be  determined  in  the  same  way  ;  but  as  the 
commercial  compounds  are  not  always  pure,  only  approximate 
results  must  be  expected. 

The  determination  of  carbon  dioxide  in  either  prepared  or 
natural  carbonates  may  be  made  with  the  apparatus  used  in 
Exp.  93.  Sodium  carbonate  and  marble  should  be  used 
for  preliminary  experiments,  as  they  are  comparatively  pure, 
and  thus  furnish  a  check  on  the  accuracy  of  the  deter- 
minations. 

CARBON  DIOXIDE  (C02). 

It  will  be  useful  at  this  point  to  examine  this  important 
compound  a  little  more  closely  than  has  already  been  done. 
It  is  a  constituent  of  the  atmosphere,  but  the  quantity  present 
varies  somewhat  under  different  conditions.  Ten  thousand 
volumes  of  air  contain  about  5  volumes  of  carbon  dioxide. 
It  is  being  constantly  passed  into  the  air  as  one  of  the 
products  of  combustion,  respiration,  and  putrefaction ;  and  as 
constantly  removed  by  growing  plants,  the  leaves  of  which 
absorb  it  readily.  It  is  then  decomposed  under  the  influence 
of  sunlight,  the  carbon  is  assimilated  by  the  growing  tissue, 
and  the  oxygen  is  returned  to  the  atmosphere.  In  this  way 
a  kind  of  rough  balance  is  struck,  and  the  average  quantity 
of  carbon  dioxide  in  the  air  kept  about  the  same. 

Preparation  of  Carbon  Dioxide.  —  If  a  considerable 
volume  of  this  gas  is  required  to  demonstrate  its  general 
properties,  it  is  most  readily  obtained  by  dissolving  calcium 


CARBON  AND  ITS  COMPOUNDS 


173 


carbonate   in   dilute   hydrochloric  acid.     The  reaction  is  as 
follows  : 

2HC1  +  CaC03  =  CaCl2  +  H20  +  C02. 

EXP.  94. — Put  some  small  pieces  of  clean  marble  into  the  bottle 
A  (Fig.  31),  and  pour  in  a  little  water  through  the  thistle  funnel  B, 
so  as  to  well  cover  the  solid.  Arrange  the 
delivery-tube  in  the  cylinder  C,  and  then 
add  strong  hydrochloric  acid,  a  little  at  a 
time,  through  B.  A  steady  stream  of  gas 
is  thus  obtained,  which,  passing  into  C, 
gradually  displaces  the  air  and  fills  the 
cylinder.  It  is  full  when  the  flame  of  a 
burning  taper  is  immediately  extinguished 
on  being  brought  into  the  mouth  of  the 
jar.  The  gas  may  also  be  collected  over 
water. 


FIG.  31. 


Properties  of  Carbon  Dioxide.— 

It  is  a  colourless,  transparent,  odourless 
gas,  and  is  somewhat  soluble  in  water, 
which  dissolves  about  1*5  times  its  own  volume  of  the  gas  at 
ordinary  temperatures.  The  solution  has  a  sour  taste,  and 
turns  litmus  red,  thus  showing  its  acid  properties.  It  is 
non-combustible,  and  a  non-supporter  of  combustion  in  the 
ordinary  way;  but  burning  magnesium  continues  to  burn 
in  the  gas  with  formation  of  magnesia  and  liberation  of 
carbon.  Ked-hot  charcoal  also  abstracts  oxygen  from  it. 
(See  Exp.  96.) 

Lime-water  is  turned  milky  when  carbon  dioxide  is  shaken 
up  with  it.  This  is  caused  by  the  formation  of  insoluble 
calcium  carbonate.  On  further  shaking  to  make  the  water 
take  up  more  gas  the  milkiness  disappears,  which  is  due  to 
the  conversion  of  the  normal  into  the  soluble  acid  carbonate. 
On  boiling  the  solution  the  normal  carbonate  is  thrown  down 
again.  The  reactions  are  shown  by  the  equations  : 

CaH202  +  C02  =  CaC03  +  H20. 
CaC03  +  C02  +  H20  =  CaH2(C03)2. 

EXP.  95. — Shake  up  some  lime-water  in  a  jar  of  carbon  dioxide 


174  METALLURGICAL  CHEMISTRY 

until  the  solution,  which  is  first  turned  milky,  becomes  clear ;  pour 
the  clear  liquid  into  a  test-tube,  and  boil  it. 

Carbon  dioxide  is  heavy,  and  can  be  poured  from  one 
vessel  to  another  in  the  ordinary  way.  Its  presence  in  the 
second  vessel  is  easily  proved  by  putting  in  a  lighted  taper, 
or  by  the  addition  of  a  little  lime-water.  It  is  22  times 
heavier  than  hydrogen,  and  1*5  times  heavier  than  air. 

Carbonic  Oxide,  CO. — The  complete  combustion  of  carbon 
in  air  results  in  the  formation  of  carbon  dioxide ;  but  if  the 
supply  of  air  is  limited,  and  the  temperature  of  the  carbon  is 
sufficiently  high,  another  oxide  of  carbon  is  produced.  It  is 
probable  that  C02  is  formed  first,  and  then  converted  into  CO 
by  abstraction  of  part  of  its  oxygen.  Thus  : 

C02  +  C=2CO. 

A  minimum  temperature  of  about  800°  C.,  or  a  bright  red 
heat,  is  required  for  the  change  to  go  on  rapidly.  A  good 
table  furnace  is  indispensable.  The  one  shown  in  Fig.  8 
has  three  Bunsen  jets  served  from  one  gas  tap,  and  the  body 
consists  of  a  rectangular  fire-brick  chamber,  surmounted  by 
a  sheet  iron  tube  to  assist  the  draught.  The  reaction-tube 
is  a  piece  of  solid-drawn  steel  tube,  such  as  is  used  for  cycle 
work,  and  about  18  inches  long.  Rubber  bungs  carry  the 
inlet-tube  A  arid  the  outlet-tube  B,  and  are  kept  cool  by  pads 
of  blotting-paper,  which  are  wrapped  round  the  ends  of  the 
steel  tube  and  kept  well  saturated  with  water.  The  charcoal 
used  should  be  previously  heated  in  a  closed  clay  crucible  to 
the  highest  temperature  of  a  wind  furnace  to  expel  hydrogen 
and  gaseous  hydrocarbons,  which  are  not  removed  during  the 
ordinary  process  of  manufacture.  This  preliminary  heating 
may  be  omitted,  but  then  the  first  two  or  three  jars  of  gas 
collected  will  be  impure. 

EXP.  96. — Insert  a  loose  plug  of  asbestos  in  the  steel  tube  about 
j  its  length  from  one  end,  and  fill  the  middle  part  with  small  pieces 
of  charcoal  about  the  size  of  a  pea.  Connect  the  apparatus  for 
generating  dry  CO.,,  A  and  B,  Fig.  36,  with  A,  Fig.  8,  and  by  the 
careful  addition  of  dilute  nitric  acid  keep  up  a  moderate  stream  of 


AND  ITS  COMPOUNDS  175 

hubbies  through  the  wash-bottle,  so  as  to  insure  a  regular  stream  of 
gas  through  the  red-hot  tube.  If  the  carbonic  oxide  is  to  be  quite 
free  from  carbon  dioxide,  a  wash-bottle  containing  a  strong  solution 
of  caustic  soda  must  be  inserted  between  the  end  of  the  steel  tube 
and  the  collecting  cylinder.  Collect  several  jars  of  the  gas. 

Properties  Of  CO.  —  It  is  a  colourless  gas  with  a  faint 
oppressive  odour,  but  must  not  be  inhaled,  as  it  is  very 
poisonous,  and  should  not  be  allowed  to  escape  unburnt.  It 
burns  with  a  pale  blue  flame,  and  explodes  when  mixed  with 
air  or  oxygen.  The  pure  gas  does  not  turn  lime-water  milky, 
but  when  it  is  burnt  the  product  of  the  combustion  does  turn 
the  lime  water  milky,  thus  proving  that  carbon  dioxide  is 
formed. 

=  2C0 


2. 

4  vols.     2  vols.     4  vols. 

CO  is  only  slightly  soluble  in  water.  Its  density,  compared 
with  air=  1,  is  0'97. 

Experiments  With  CO.  —  (1)  Add  a  little  lime-water  to  a 
jar  of  the  pure  gas,  and  shake  the  jar;  burn  the  gas  and 
shake  the  jar  again.  (2)  Mix  1  volume  of  the  gas  with  2'5 
volumes  of  air  in  a  tall  cylinder,  and  ignite  the  mixture.  The 
flame  runs  rapidly  down  the  jar  and  thus  furnishes  an  example 
of  incipient  explosion.  (3)  Mix  2  volumes  of  the  gas  with 
1  volume  of  oxygen,  and  ignite  the  mixture.  The  combustion 
is  much  more  rapid,  and  a  decided  explosion  is  the  result. 
The  violence  of  the  explosion  when  an  explosive  mixture  of 
gases  is  ignited,  depends  upon  the  rapidity  with  which  the 
combustion  extends  through  the  mass. 

Formation  of  CO  when  a  Limited  Supply  of  Air  is  passed 
through  Red-hot  Carbon.  EXP.  97.  —  Disconnect  the  CO2  apparatus 
from  the  tube  used  in  the  last  experiment,  and  replace  it  by  any 
convenient  apparatus  for  sending  a  slow  stream  of  dry  air  through 
the  red-hot  charcoal.  The  form  shown  in  Fig.  32  is  very  simple.  B  is 
full  of  water  ;  C  is  full  of  air  ;  D  contains  some  strong  sulphuric  acid. 
By  running  the  water  slowly  from  B  into  C  a  slow  stream  of  dry 
air  is  forced  through  the  red-hot  charcoal,  when  the  drying-bottle 
D  is  connected  with  the  tube  containing  the  charcoal.  If  the 
issuing  gas  is  required  to  be  free  from  CO2,  the  traces  of  this  gas 


176 


METALLURGICAL  CHEMISTRY 


B 


can  be  removed  by  caustic  soda,  as  described  above.    Collect  two  or 
three  jars  of  the  gas.     Add  some  lime-water  to  a  jar  of  the  gas 

which  has  been  puri- 
fied by  caustic  soda. 
If  the  liquid  remains 
clear  the  gas  is  free 
from  CO2.  Put  a 
lighted  taper  to  the 
mouth  of  the  jar,  and 
when  the  combustion 
has  ceased,  shake  up 
the  gas  with  the 
lime  -  water,  which 
turns  milky,  thus 
proving  the  presence 
of  the  CO.,,  formed 
by  the  combustion  of 
the  CO.  Inflame 
another  jar  of  the 

DJIIJII                             I'll  |      gas,  and  notice  that 

ll_ — „ JPA      it     burns    with     the 

characteristic  blue 
flame  of  CO. 


FIG. 


The  carbonic 
oxide  formed  oc- 
cupies twice  the  volume  of  the  oxygen  absorbed,  and  is  also 
mixed  with  the  nitrogen  of  the  air  from  which  the  oxygen  is 
taken.  Thus  the  mixture  contains  one-third  CO  and  two- 
thirds  N2  by  volume. 

Action  of  Red-hot  Carbon  on  Water  Vapour.—  The 

principal  reaction  which  takes  place  is  one  in  which  carbonic 
oxide  is  formed  and  hydrogen  gas  liberated. 


Another  reaction  which  takes  place  in  a  much  more  limited 
degree  brings  about  the  formation  of  marsh  gas,  CH4,  and 
probably  other  hydrocarbons  of  a  similar  character. 


20  =  CH4-f2CO. 

If  the  temperature  falls  too  low,  hydrogen  and  carbon  dioxide 
are  formed. 


CARBON  AND  ITS  COMPOUNDS  177 

So  that  when  steam  is  passed  through  red-hot  carbon  at  least 
three  combustible  gases  are  formed.  Both  reactions  are 
endothermie— i.e.,  more  heat  is  absorbed  in  effecting  the 
decomposition  of  the  water  taking  part  in  the  reactions  than 
is  evolved  by  the  formation  of  the  resulting  compounds  when 
the  change  takes  place ;  and  there  is  a  definite  limiting  tem- 
perature below  which  it  ceases.  It  is  therefore  necessary  to 
keep  the  temperature  above  this  limit  by  the  application  of 
heat  from  outside,  or  by  allowing  some  of  the  carbon  to  burn 
to  carbon  dioxide.  Similar  remarks  apply  to  the  reaction 
C  +  C02  =  2CO,  which  is  also  endothermie.  On  the  other 
hand,  C  +  02  =  C02  is  an  exothermic  change— i.e.,  heat  is 
given  out  while  it  is  taking  place. 

EXP.  98. — Use  the  apparatus  shown  in  Fig.  8,  but  replace  the 
iron  by  charcoal.  Eegulate  the  boiling  of  the  water  in  the  flask  so 
as  to  give  a  steady  stream  of  bubbles  from  the  delivery-tube  into 
the  collecting  cylinder.  Collect  two  or  three  jars  of  the  gas.  It  is 
a  mixture  of  carbonic  oxide,  hydrogen,  and  marsh  gas,  and  is  called 
"  water  "  gas.  Burn  a  jar  of  the  gas,  and  note  the  faintly  luminous 
flame.  Mix  1  volume  of  the  gas  with  2£  volumes  of  air,  and 
explode  the  mixture.  The  explosion  is  louder  than  ..'th  carbonic 
oxide  alone. 

Ordinary  air  contains  water  vapour,  and  if  it  is  used  in 
Exp.  97  instead  of  dry  air,  a  small  amount  of  hydrogen  is 
present  in  the  gaseous  mixture  obtained. 

Coal  is  a  mineral  of  vegetable  origin  and  of  very  complex 
.composition.  It  contains  C,  H,  0,  N,  and  ash-forming  con- 
stituents. Sulphur  is  also  invariably  present  in  small  quantity. 
The  whole  of  the  carbon  and  part  of  the  hydrogen  are  avail, 
able  for  the  development  of  heat  when  the  coal  is  burnt.  In 
practice  it  is  usual  to  subtract  from  the  total  hydrogen  one- 
eighth  the  weight  of  the  oxygen  present,  and  to  consider  that 
quantity  of  hydrogen  as  already  combined  in  the  form  of 
water,  and  therefore  not  available  for  heat  production.  The 
remainder  is  known  as  available  hydrogen.  The  older  the 
coal  the  more  concentrated  it  is  ;  for  the  changes  which  the 
vegetable  matter  undergoes  during  its  extremely  slow  conver- 

12 


178 


METALLUKGICAL  CHEMISTKY 


sion  into  coal  cause  the  percentage  of  carbon,  and  also  the 
ratio  of  the  hydrogen  to  the  oxygen,  to  increase.  That  is,  the 
percentages  of  oxygen  and  hydrogen  are  both  decreased,  but 
the  former  more  than  the  latter. 

Coal  containing  much  hydrogen  undergoes  very  complicated 
changes  when  it  is  strongly  heated  out  of  contact  with  air. 
The  greater  part  of  the  oxygen  is  liberated  as  water;  the 
nitrogen  principally  in  the  form  of  ammonia,  and  the  residue 
of  the  hydrogen  partly  in  the  free  state,  and  partly  in  a  very 
numerous  and  complex  series  of  carbon  compounds. 

EXP.  99. — Draw  off  a  piece  of  combustion-tube  in  the  blow-pipe 
flame  ;  put  in  it  a  few  grams  of  powdered  coal ;  draw  off  and  bend 
the  thinned-out  portion,  as  shown  in  Fig.  33.  A  is  the  loaded  tube 
in  position ;  B  is  a  bottle  to  condense  liquid  matter ;  C  is  a  cylinder 
to  collect  the  escaping  gas.  Strongly  heat  the  tube  as  long  as  gas 


FIG.  33. 

collects  in  C.  Remove  the  cylinder,  and  put  a  light  to  the  mouth. 
The  gas  burns  with  a  bright  white  flame.  B  is  found  to  contain  a 
watery  liquid  and  a  black  oily  liquid.  On  breaking  A,  a  greyish- 
black  coherent  mass  of  coke  is  obtained. 


The  coal  is  split  up  into:  (1)  a  combustible  gas;  (2)  a 
watery  liquid  which  contains  ammonia;  (3)  tar,  a  liquid  of 
very  complex  character;  (4)  coke,  a  solid  residue  containing 
carbon  (about  90  per  cent.),  some  hydrogen  and  oxygen,  and 
inorganic  matter  or  ash. 


CAEBON  AND  ITS  COMPOUNDS  179 

APPROXIMATE  COMPOSITION  OF  COAL  GAS. 

Hydrogen,  H2     ...  ...  ...  ...  45 

Marsh  gas,  CH4  ...  ....  ...  ...  40 

Carbonic  Oxide,  CO  ...  ...  ...  5 

Hydrocarbons      ...  ...  ...  ...  5 

Non-combustible  gases  ...  ...  ...  5 

100 

The  experiment  just  described  illustrates  in  a  general  way 
the  principles  involved  in  the  manufacture  of  illuminating 
gas.  If  wood  is  substituted  for  coal  in  Exp.  99,  a  similar 
result  is  obtained,  and  so  striking  is  the  analogy  that  one  is 
forced  to  admit  the  similarity  in  the  composition  of  wood  and 
coal,  and  the  probability  that  they  have  the  same  origin. 

The  Changes  in  an  Open  Coal  Fire.— When  fresh  coal 

is  put  on  an  open  fire  the  heat  from  the  combustion  of  the 
fuel  already  in  the  grate  brings  about  the  changes  described 
above.  Water  is  expelled,  and  forms  the  greater  part  of  the 
first  smoke.  Then  the  volatile  matter  becomes  combustible, 
and  burns  with  a  white  flame  in  the  current  of  air  passing 
over  the  fire  on  its  way  to  the  chimney.  The  coal  is  gradually 
coked,  and  the  fire  settles  down  to  steady  combustion. 
Another  current  of  air  passing  through  the  fire  by  way  of 
the  grate-bars  gives  up  its  oxygen  to  the  incandescent  carbon 
as  soon  as  it  enters,  and  carbon  dioxide  is  formed.  This  gas 
is  carried  into  the  heart  of  the  fire,  there  to  be  converted  into 
carbonic  oxide  by  the  red-hot  carbon.  On  emerging  from  the 
upper  surface  of  the  fire  this  carbonic  oxide  burns  with  its 
characteristic  blue  flame  in  the  stream  of  air  passing  over  the 
top.  The  water  vapour  in  the  lower  current  of  air  also  under- 
goes similar  changes  to  those  already  described,  hydrogen  and 
carbonic  oxide  being  added  to  the  combustible  gas.  The 
earthy  matter  of  the  coal  is  left  in  the  grate  as  ash.  It  is 
evident,  from  the  accumulation  of  soot  in  the  chimney,  and 

12—2 


180  METALLURGICAL  CHEMISTRY 

the  contamination  of  the  atmosphere,  that  combustion  in  an 
open  grate  is  far  from  perfect. 

Closed  Grate. — If  the  top  of  the  grate  is  closed  so  as  to 
cut  off  the  upper  air  current,  and  the  bar  spaces  are  arranged 
to  admit  a  limited  supply  of  air,  the  combustion  is  modified, 
and  the  gas  drawn  off  from  the  closed  top  contains  upwards 
of  30  per  cent,  of  combustible  matter.  This  is  known  as 
*  air '  gas  or  '  producer '  gas,  and  is  made  on  the  large  scale  for 
heating  purposes.  By  introducing  more  water  vapour,  up  to 
10  per  cent.,  gas  containing  a  larger  percentage  of  hydrogen 
is  obtained.  The  most  recent  improvements  are  in  this 
direction. 

AVERAGE  COMPOSITION  OF  '  Am '  GAS. 

Carbonic  oxide,  CO  ...  ...  ...  25 

Carbon  dioxide,  C02  ...  ...  ...  5 

Hydrogen,  H2     ...  ...  ...  ...  10 

Nitrogen,  N2       ...  ...  ...  ...  55 

Hydrocarbons      ...  ...  ...  ...  5 

100 

The  changes  are  for  the  most  part  endothermic,  and  the 
complete  combustion  of  a  small  percentage  of  the  carbon  must 
take  place  to  supply  the  necessary  heat  for  the  endothermic 
changes. 

SUMMARY. 

Carbon  is  found  in  the  free  state  in  Nature.  It  is  also  a 
constituent  of  many  bodies  of  very  diverse  composition.  It 
exists  in  three  distinct  allotropic  modifications.  The  car- 
bonates are  a  very  important  class  of  metallic  compounds,  and 
some  of  them  are  used  for  the  extraction  of  the  metals  they 
contain.  They  are  readily  decomposed  by  heat  into  metallic 
oxides  and  carbon  dioxide.  Acids  dissolve  them  with  evolu- 
tion of  the  same  gas.  There  are  only  two  oxides  of  carbon. 
The  dioxide  is  an  acid-forming  oxide,  and  the  monoxide  is  a 


CARBON  AND  ITS  COMPOUNDS  1S1 

very  useful  combustible  gas.  Coal  and  its  derivatives,  coke 
and  combustible  gas,  are  extensively  used  in  the  extraction 
and  after-treatment  of  the  common  metals. 

QUESTIONS. 

1.  Give  a  short  account  of  the  different  forms  of  carbon. 
How  are  they  proved  to  be  one  and  the  same  element  ? 

2.  What  is  the  action  of  heat  on  carbonates  generally  ? 

3.  What  are  the  properties  of  carbon  dioxide,  and  how  may 
it  be  converted  into  carbon  monoxide  ? 

4.  What  happens  when  coal  or  wood  is  heated  in  a  closed 
vessel  ? 

5.  Explain  fully  what  takes  place  when  steam  is  brought 
into  contact  with  red-hot  carbon. 

6.  Describe  the  changes  taking  place  in  an  ordinary  coal 
fire.     What   is  the  result   of   closing  the  top  of  the  grate  ? 
Give  the  equations  expressing  the  various  changes. 


CHAPTER  XI 
EEDUCTION 

REDUCTION  is  a  term  used  by  metallurgists  to  express  the 
chemical  change  which  takes  place  when  a  metal  is  set  free 
from  its  combinations  with  non-metals.  As  oxides  and 
sulphides  are  the  most  abundant  natural  compounds  of  the 
metals,  they  furnish  frequent  examples  of  reducing  action, 
and  are  the  chief  sources  of  the  prepared  metals.  But  for 
illustrating  general  principles  it  will  be  more  convenient  to 
use  prepared  compounds,  and  so  avoid  the  complications 
which  would  be  introduced  by  the  presence  of  impurities  in 
the  natural  materials. 

Reduction  by  Heat  alone.  The  simplest  change  is  that 
which  takes  place  when  an  oxide  is  decomposed  by  heat. 
There  are  very  few  examples  of  this  kind  of  change,  for  the 
oxides  generally  are  stable  bodies,  and  not  readily  decomposed 
by  heat  alone.  Oxides  of  mercury,  silver,  and  gold  are  the 
only  common  examples.  The  decomposition  of  red  oxide  of 
mercury  (p.  39)  may  be  taken  as  typical  of  the  change.  But 
when  the  compounds  are  salts  of  these  oxides  the  decomposi- 
tion is  more  complex.  Thus  with  silver  nitrate,  nitrogen 
peroxide  and  oxygen  are  set  free  along  with  the  metal. 

2AgN03=  2Ag  +  2N02  +  02. 

EXP.  100. — Put  a  crystal  of  silver  nitrate  into  a  small  dry  test- 
tube,  and  heat  it  in  the  Bunsen  flame.  Note  the  red  gas  which  is 
given  off,  and  test  it  with  a  glowing  splint.  Remove  the  reduced 
silver  from  the  tube,  and  rub  it  on  a  hard  surface  with  the  blade  of 
a  knife.  The  particles  of  the  reduced  metal  are  pressed  together, 

182 


REDUCTION  183 

and  present  the  characteristic  metallic  appearance.  The  experiment 
may  be  made  quantitative  by  weighing  the  nitrate  (about  0'5  gram) 
before,  and  the  residue  of  silver  after,  heating. 

A  partial  reduction  is  often  effected  by  heat  alone,  and  a 
simpler  compound  of  the  metal  thus  obtained.  This  is  notably 
so  with  the  metallic  carbonates,  of  which  some  examples  have 
been  given  (p.  169).  The  decomposition  of  lead  carbonate  is 
easily  brought  about,  but  the  residue  of  lead  oxide  is  some- 
what readily  fused,  and,  if  allowed  to  get  too  hot,  would  spoil 
a  porcelain  crucible.  Both  the  normal  and  the  basic  carbonates 
are  prepared,  and  a  simple  experiment  will  distinguish  between 
them. 

EXP.  101. — Weigh  a  convenient  piece  of  a  broken  porcelain  dish, 
and  weigh  on  it  1  gram  of  lead  carbonate.  Put  the  porcelain  on  a 
pipeclay  triangle  and  heat  it  until  partial  fusion  takes  place.  Allow 
the  porcelain  to  cool,  and  re  weigh  it. 

The  loss  is  due  to  the  escape  of  carbon  dioxide,  and  the 
result  may  be  compared  with  the  equation — 

PbC03  =  PbO  +  C02. 

A  partial  reduction  also  takes  place  when  some  sulphates  are 
strongly  heated,  but  the  reaction  is  not  quite  so  general  as 
with  the  carbonates.  (See  Exp.  86.) 

EXP.  102. — Weigh  2  grams  of  powdered  green  vitriol  in  a  weighed 
porcelain  crucible,  and  heat  it  over  a  small  Bunsen  flame  for  fifteen 
minutes.  Allow  the  crucible  to  cool,  and  reweigh  it.  The  loss  is 
due  to  the  escape  of  the  water  of  crystallization.  Then  put  the 
crucible  into  a  gas  muffle,  and  keep  it  at  a  moderate  red  heat  for 
thirty  minutes.  When  the  crucible  is  cold,  reweigh  it.  The  second 
loss  is  due  to  the  escape  of  oxides  of  sulphur.  The  residue,  if 
not  overheated,  is  a  red  powder.  Compare  the  results  with  the 
equation : 

2(FeS04,  7H20)=Fe203+S02+S03+14H20. 

Residue.     Second  loss.     First  loss. 

This  experiment  illustrates  the  method  used  in  the  preparation 
of  rouge,  and  as  the  colour  depends  largely  on  the  temperature 
at  which  the  reduction  takes  place,  care  must  be  exercised  in 


184  METALLURGICAL  CHEMISTRY 

carrying  it  out.  The  value  of  the  rouge  also  depends  largely 
on  the  grinding  and  levigation  which  the  residue  undergoes 
after  the  reduction.  (See  Exp.  3.) 

The  partial  reduction  of  metallic  sulphides  rich  in  sulphur 
is  demonstrated  in  Exp.  38.  And  it  may  be  taken  generally 
that  when  there  are  two  or  more  sulphides  of  the  same  metal, 
the  one  containing  the  smallest  proportion  of  sulphur  is  the 
most  stable,  and  will  withstand  a  high  temperature  without 
decomposition.  The  others  are  reduced  to  this  stable  form 
when  strongly  heated.  Iron  and  copper  sulphides  furnish 
good  examples  : 

1. 

2. 

The  lower  sulphide  in  each  case  may  be  fused  at  a  bright  red 
heat  out  of  contact  with  air. 

Reduction  by  Electricity.  —  The  decomposition  of  water 
(p.  27)  is  a  simple  reduction  process  as  far  as  the  observed 
effects  are  concerned.  For,  though  a  little  sulphuric  acid  is 
necessary  to  the  success  of  the  experiment,  the  final  result 
is  expressed  thus  : 


If  a  solution  of  copper  sulphate  is  substituted  for  the  acid 
solution  in  the  voltameter,  a  deposit  of  copper  is  formed  on 
the  plate  from  which  the  hydrogen  escapes  during  the 
decomposition  of  water,  and  oxygen  gas  escapes  from  the 
other  plate.  The  final  result  is  therefore  equivalent  to  the 
reduction  of  copper  oxide,  whatever  the  intermediate  changes 
may  be. 

2CuO  =  2Cu  +  02. 

Free  sulphuric  acid  appears  in  the  solution.  This  is  probably 
due  to  secondary  changes  in  which  water  takes  part,  and 
furnishes  the  evolved  oxygen. 

1.  CuS04  =  Cu  +  (S04) 

Cathion.   Anion. 

2. 


REDUCTION  185 

Similar  reactions  take  place  when  a  current  of  electricity  is 
passed  through  a  solution  of  any  metallic  salt,  and  if  the 
liberated  metal  does  not  act  upon  water,  it  is  deposited  upon 
the  cathode  plate  (see  p.  107).  But  if  the  metal  reacts  with 
water  its  hydroxide  is  formed,  and  hydrogen  is  liberated. 
Common  salt  fuses  to  a  clear  liquid  at  a  red  heat,  and  if  a 
current  of  electricity  is  passed  through  the  molten  mass 
between  carbon  electrodes,  the  metal  sodium  and  the  gas 
chlorine  are  liberated  and  can  be  readily  collected. 


Sodium 
chloride. 


Similarly,  the  cheap  production  of  aluminium  depends  upon 
the  passage  of  an  electric  current  through  fused  cryolite, 
AlF3,3NaF,  in  which  alumina  is  dissolved.  The  cryolite  is 
unchanged. 


2A1203  =  4 

Reduction  by  Heat  and  Carbon.  —Every  compound 

tends  to  decompose  when  it  is  heated,  and  the  higher  the 
temperature  to  which  it  is  exposed^  the  greater  its  proneness 
to  decomposition.  Many  compounds,  however,  resist  a  high 
temperature,  and  if  their  reduction  is  to  be  effected  the  action 
of  heat  must  be  assisted  by  that  of  a  reducing  agent,  which 
may  be  either  an  element  or  another  compound.  The  reduc- 
ing agent  exerts  a  pull  upon  one  or  more  constituents  of  the 
compound,  and  thus  assists  the  heat  to  effect  its  decomposition. 
The  action  is  most  effective  when  the  reducing  agent  forms  a 
compound  which  is  volatile  at  the  temperature  of  the  reduc- 
tion, and  is  rapidly  removed  from  the  sphere  of  action.  But 
this  is  not  a  necessary  condition.  Carbon,  when  it  helps  to 
reduce  a  metallic  oxide,  forms  either  carbon  dioxide  or  carbon 
monoxide,  according  to  the  temperature  necessary  to  bring 
about  the  reduction. 

EXP.  103.  —  Well  mix  0'2  gram  of  powdered  charcoal  with  8  grams 
of  lead  oxide  ;  put  the  mixture  into  a  small  test-tube,  and  tit  the 


186 


METALLUEGICAL  CHEMISTEY 


tube  with  a  bung  and  delivery-tube,  as  shown  in  Fig.  34.     Heat 
the    mixture   strongly  for   five  minutes,  and  collect  the  escaping 

gas  in  the  jar.  Test  the  gas 
with  a  lighted  taper ;  it  is  a  non- 
supporter  of  combustion.  Pour 
the  gas  from  the  jar  into  a 
beaker  containing  a  little  clear 
lime  -  water,  and  shake  the 
beaker ;  the  gas  is  heavy,  and 
turns  lime-water  milky.  It  is 
carbon  dioxide,  C02.  Examine 
the  residue  in  the  tube  for  shots 
of  metallic  lead. 


A  similar  experiment  may 
be  made  with  black  oxide  of 
but    the    proportion 


FIG.  34. 

copper 

between  the  charcoal  and  the   oxide  must  be  nearly  three 
times  as  great. 


EXP.  104. — Put  about  2  grams  of  powdered  wood  charcoal  into  a 
porcelain  or  platinum  crucible ;  close  the  crucible  with  a  lid  to 
prevent  the  admission  of  air,  and  heat  it  strongly  over  the  blow-pipe 
flame  for  ten  minutes.  Allow  the  crucible  to  get  cold  before 
removing  the  lid.  Charcoal  thus  strongly  heated  contains  about 
92  per  cent,  of  carbon.  Weigh  0  054  gram  of  this  charcoal,  which 
is  roughly  equal  to  0'05  gram  of  carbon.  Mix  it  thoroughly  with 
4  grams  of  dry  litharge,  and  introduce  the  mixture,  with  the  help 
of  a  paper  gutter,  into  a  piece  of  combustion  tube,  ^  inch  by 
4  inches,  closed  at  one  end.  Weigh  the  tube  and  mixture  together ; 
fix  the  tube  in  a  clip,  and  heat  it  strongly  in  the  Bunsen  flame 
until  the  mixture  becomes  pasty.  Allow  the  tube  to  cool,  and 
reweigh  it. 

The  carbon  is  converted  into  carbon  dioxide  by  oxygen 
obtained  from  the  lead  oxide,  and  the  excess  of  the  oxide 
insures  the  combustion  of  the  whole  of  the  carbon. 

EXAMPLE. — Weight  of  tube  +  mixture  =     18 '983  grams 

„  „        after  heating     =     18*799      „ 

Loss         =      0184  gram 

The  loss  is  due  to  the  escape  of  carbon  dioxide  containing 
0'05  gram  of  carbon.  Therefore  the  weight  of  oxygen  com- 


KEDUCTION 


187 


bined  with  this  weight  of  carbon  is  (H34  gram.   These  weights 
may  be  used  to  find  the  formula  of  the  compound. 


=  0-0083.     .-.  The  ratio  is  1:2, 


For  ^  =  0-0041;  and  ^| 

and  the  formula  is  C02, 

Compare  the  results  of  the  experiment  with  the  equation — 


EXP.  105. — Draw  off  a  piece  of  £-inch  combustion-tube  as  shown 
in  Fig.  85.  Weigh  accurately  3  grams  of  lead  oxide,  and  transfer 
it  to  the  tube  with  the 
help  of  a  paper  gutter. 
Weigh  the  tube  and  oxide 
together.  Fix  it  in  a 
clip,  and  connect  the  thin 
end  with  a  rubber  tube  to 
the  gas -tap.  Turn  on  the 
gas,  and  ignite  it  at  the 
open  end  of  the  tube, 
allowing  it  to  burn  with 
a  moderate  flame .  Heat 
the  tube  in  the  Bunsen 
flame,  and  watch  the  re- 
duction as  it  proceeds. 
Note  that  moisture  col-  FIG.  35. 

lects  on  the  cool  part  of 

the  tube  at  first,  but  disappears  as  it  gets  warm.  When  the  reduc- 
tion is  complete,  which  is  indicated  by  the  disappearance  of  the 
oxide,  remove  the  burner,  but  allow  the  gas  to  continue  passing 
through  the  tube  until  it  has  cooled  somewhat ;  then  reweigh  it. 

Coal  gas  contains  carbon  and   hydrogen,    and   the   water 
formed  is  the  result  of  the  oxidation  of  the  latter  (p.  189). 

EXAMPLE. — 3   grams  of  lead  oxide  lost  0*219  gram  of  oxygen. 


Then 


0-219 


=  0-0137;  and 


=  0  0134,  so  that  the  ratio  is  1:1. 


16  207 

Therefore  the  formula  of  the  oxide  used  is  PbO. 

An  exactly  similar  experiment  can  be  made  with  copper 
oxide. 

The  temperature  required  for  the  above  reduction  is  com- 
paratively low,  and  carbon  dioxide  is  formed ;  but  when  a 


188  METALLUEGICAL  CHEMISTRY 

high  temperature  is  necessary  for  the  reduction  carbonic 
oxide,  CO,  is  the  result.  This  is  not  easily  demonstrated, 
and  must  be  left  for  more  advanced  work.  The  reduction  of 
tin  and  zinc  oxides  may  be  cited  as  examples,  and  the 
simplest  way  of  representing  the  reactions  is  shown  in  the 
equations — 

Sn02+2C  =  Sn  +  2CO. 
ZnO  +  C  =  Zn  +  CO. 

But  it  is  probable  that  carbon  dioxide  is  first  formed,  and 
then  reduced  by  excess  of  carbon  at  the  high  temperature 
required.  Some  oxides,  such  as  alumina  and  lime,  require  the 
excessively  high  temperature  of  the  electric  furnace  to  effect 
their  reduction. 

Reduction  by  Hydrogen.— The  action  of  hydrogen  gas 
upon  hot  copper  oxide  has  received  more  attention  from 
eminent  chemists  than  any  other  experiment  in  the  whole 
range  of  the  science.  It  was  used  by  Dumas  in  his  deter- 
mination of  the  combining  proportion  of  oxygen,  and  his 
experiment  has  been  repeated  several  times,  with  every  possible 
refinement  of  experimental  research.  For  full  details  of  this 
work  some  standard  book  on  the  subject  must  be  consulted ; 
but  a  comparatively  simple  experiment  may  be  described,  and 
even  carried  out  with  ordinary  care.  In  this  case  approximate 
values  only  must  be  expected. 

EXPERIMENT  TO  PROVE  THE  COMPOSITION  OF  WATER  BY 

WEIGHT. 

The  apparatus  shown  in  Fig.  36  is  to  be  fitted  up.  A  is  a 
bottle  fitted  with  a  tap  funnel,  and  used  for  preparing  the 
hydrogen.  B  is  a  wash-bottle  containing  a  little  water  to 
wash  the  gas.  C  is  a  drying-tube  filled  with  fused  calcium 
chloride  and  glass-wool.  D  is  a  combustion-tube  containing 
ignited  copper  oxide.  E  is  the  water-absorbing  portion,  which 
consists^of  a  test-tube,  the  upper  part  of  which  is  filled  with 


EEDUCTION 


189 


glass-wool,  and  a  calcium  chloride  drying-tube.  The  delivery- 
tube  from  C  dips  into  a  small  beaker  containing  a  little  strong 
sulphuric  acid.  This  serves  to  show  the  rate  at  which  the 
excess  of  hydrogen  is  escaping,  and  also  prevents  the  absorption 
of  moisture  from  the  air. 

To  make  the  experiment,  D  and  E  are  weighed  separately, 
the  rubber  joints  a,  6,  and  c  being  closed  by  the  insertion  of 
bits  of  glass  rod.  The  connections  are  then  made,  and  a 
moderate  current  of  hydrogen  passed  through  the  apparatus, 
escaping  through  the  acid  in  the  beaker.  When  the  air  has 
been  displaced,  the  copper  oxide  in  D  is  heated ;  the  reduction 
commences,  and  is  continued  until  a  quantity  of  water  has 


FIG 


collected  in  the  test-tube  of  E.  The  flame  is  then  removed, 
and  the  tube  allowed  to  cool  with  the  hydrogen  still  passing 
through  it.  The  parts  are  then  separated  and  reweighed.  If 
the  experiment  is  conducted  with  care,  it  proves  approximately 
that  the  loss  of  weight  in  D  is  to  the  gain  of  weight  in  E  as 
8  is  to  9 ;  or  that  8  parts  by  weight  of  oxygen  combine  with 
1  part  by  weight  of  hydrogen  to  form  9  parts  by  weight  of 
water.  The  reduction  is  expressed  by  the  equation — 

CuO+H2  =  Cu+H20. 

Ferric  oxide,  Fe203,  is  also  readily  reduced  by  hydrogen,  and 
the  reaction  may  be  made  the  subject  of  a  simple  experiment. 


190  METALLURGICAL  CHEMISTRY 

EXP.  106.  —  Fit  up  the  apparatus  shown  in  Fig.  36,  with  the 
exception  of  the  part  E.  Put  2  grams  of  ferric  oxide  into  the  tube  D, 
and  pass  a  steady  current  of  hydrogen  through  it.  When  the  air 
has  been  displaced,  -heat  the  tube  until  the  red  powder  turns  com- 
pletely black.  Allow  the  tube  to  get  cold  while  the  gas  is  still 
passing  through  it  ;  then  transfer  the  reduced  metal  to  a  dry  test- 
tube  and  cork  it  up,  to  exclude  the  air  from  the  metal. 

The  condensation  of  water  in  the  cold  neck  of  the  com- 
bustion-tube is  readily  observed.  The  equation  is  : 


If  ferric  oxide  is  reduced  at  as  low  a  temperature  as  possible, 
a  black  pyrophoric  powder  is  obtained,  which  takes  fire  spon- 
taneously if  allowed  to  come  into  contact  with  the  air  ;  but  if 
reduced  at  a  higher  temperature  the  metal  is  much  more 
stable,  and  may  be  exposed  to  the  air  when  cold  without 
alteration.  If  a  gram  of  this  reduced  metal  is  heated  in  an 
open  crucible,  it  increases  to  1*43  grams,  which  corresponds 
to  the  formula  Fe2O3.  The  iron  is  again  converted  completely 
into  its  oxide. 

A  number  of  other  metallic  oxides  are  reduced  by  hydrogen 
under  similar  conditions. 

Reduction  by  Carbonic  Oxide.  —  The  action  of  carbonic 

oxide  is  almost  as  general  as  that  of  carbon  and  hydrogen. 
It  may  therefore  be  described  as  a  common  reducing  agent 
for  oxides.  The  reduction  of  oxide  of  iron  is  very  interesting, 
because  the  reduced  metal  reacts  on  the  excess  of  the  reducing 
gas  and  absorbs  carbon  from  it.  The  first  change  is  definite, 
and  is  expressed  thus  : 


It  is  found  that  a  considerable  excess  of  carbonic  oxide  must 
be  used,  or  the  reduction  is  not  complete.  The  reaction  is  the 
principal  reducing  one  taking  place  in  an  iron-smelting  blast 
furnace,  and  it  is  well  known  that  there  must  be  at  least  twice 
as  much  carbonic  oxide  as  carbon  dioxide  in  the  waste  gas  if 
the  furnace  is  to  work  satisfactorily.  The  general  reactions 


REDUCTION  191 

by  which  the  reducing  gas  is  formed  in  the  blast  furnace  are 
described  on  p.  174;  but  by  using  a  forced  draught  a  very 
much  higher  temperature  is  obtained  in  the  lower  part  of  the 
furnace  than  in  a  closed  grate.  For  the  fusion  of  the  slag  and 
of  the  reduced  metal,  as  well  as  the  formation  of  carbonic 
oxide,  has  to  be  considered.  A  falling  off  in  the  proportion 
of  carbonic  oxide  would  result  in  a  partial  reduction  by  which 
ferrous  oxide  would  be  formed,  and  pass  into  the  slag,  with 
consequent  loss  of  iron.  The  reaction  is  : 


Ferric  Ferrous 

oxide  oxide 

The  second  change,  in  which  the  reduced  metal  reacts  with 
carbonic  oxide  and  abstracts  carbon  from  it,  is  not  so  definite 
as  the  actual  reduction  of  the  metal,  and  is  usually  considered 
generally.  Thus,  if  x  is  used  to  represent  an  unknown  but 
determinate  quantity,  the  reaction  may  be  represented  thus  : 

xFe  +  2CO  =  Fe^C  +  C02. 

The  reactions  expressed  in  the  last  three  equations  are  readily 
demonstrated,  but  the  first  and  third  only  need  be  considered 
here.  The  carbonic  oxide  required  may  be  prepared  in  any 
convenient  way;  but  the  decomposition  of  yellow  prussiate 
of  potash  by  strong  sulphuric  acid  furnishes  a  ready  means  of 
obtaining  the  pure  gas. 

EXP.  107.  —  Fit  up  a  2-litre  bottle,  as  shown  in  Fig.  37,  to  serve 
as  a  simple  gas-holder.  Put  6  grams  of  powdered  yellow  prussiate 
of  potash  into  an  8  ounce  flask,  cover  it  with  strong  sulphuric  acid, 
and  fit  the  neck  with  a  bung  and  delivery-tube.  Fill  the  gas-holder 
with  water  ;  connect  the  flask  with  the  tube  A,  and  gently  heat  it 
on  gauze  over  a  small  flame.  Carbonic  oxide  will  be  given  off 
steadily  if  the  flask  is  not  heated  too  rapidly,  and  will  displace  the 
water  through  the  tube  B.  The  quantity  of  yellow  prussiate  used 
should  be  arranged  so  that  the  liberated  gas  does  not  quite  fill  the 
gas-holder.  When  the  evolution  of  gas  ceases  close  the  tube  A  with 
a  pinch  tap,  and  the  apparatus  is  ready  for  use.  Draw  off  a  piece 
of  combustion-tube  similar  to  that  used  in  Exp.  105  ;  put  into  it 
2  grams  of  ferric  oxide,  and  arrange  it  in  a  clip  for  heating  over  the 
Bunsen  burner.  Connect  A,  Fig.  37,  with  a  drying-bottle  contain- 


192 


METALLURGICAL  CHEMISTRY 


FIG.  37. 


ing  a  little  strong  sulphuric  acid,  and  B  with  the  tap  of  an  aspirator 

filled  with  water.  Place  the  aspirator  on  a  stand  so  that  its  bottom 
is  on  a  level  with  the  tube  A.  Connect  the 
outlet  tube  of  the  drying-bottle  with  the 
drawn- off  end  of  the  reduction -tube ;  heat 
the  part  of  the  tube  containing  the  ferric 
oxide  with  a  Bunsen  flame,  and  drive  a 
steady  stream  of  carbonic  oxide  over  it  by 
opening  the  tap  of  the  aspirator.  The  flow 
of  water  through  B  forces  the  gas  through 
A,  and  the  rate  of  flow  can  be  easily  regu- 
lated. Place  a  Bunsen  flame  near  the  out- 
let of  the  reduction-tube  to  burn  the  excess 
of  carbonic  oxide  as  it  escapes.  This  is 
most  important  on  account  of  the  very 
poisonous  nature  of  the  gas.  When  the 
reduction  is  complete,  fit  a  cork  in  the  end 
of  the  tube  to  prevent  admission  of  air, 
and  turn  off  the  tap  of  the  aspirator  at  the 

same  time.     When  the  reduced  metal  is  cold  transfer  it  to  a  dry 

test-tube,  and  cork  it  up. 

If  a  weighed  quantity  of  this  reduced  metal  is  heated  in  an 
open  porcelain  crucible,  it  is  changed  into  the  red  oxide  again, 
but  the  increase  in  weight  falls  considerably  short  of  that 
obtained  in  Exp.  106  with  the  pure  metal.  This  is  due  to  the 
presence  of  carbon,  which  is  proved  conclusively  by  the 
following  experiment : 

EXP.  108. — Transfer  a  little  of  the  reduced  metal  to  a  combustion- 
tube  ;  aspirate  a  current  of  air  over  it  and  through  a  wash-bottle 
containing  lime-water;  heat  the  part  of  the  tube  containing  the 
metal,  and  note  the  milky  precipitate  formed  in  the  lime-water. 
This  proves  without  doubt  that  carbon  is  associated  with  the  iron 
when  it  is  reduced  by  carbonic  oxide. 

Coal  gas  contains  about  5  per  cent,  of  carbonic  oxide,  arid 
if  used  for  the  reduction  of  ferric  oxide  the  reduced  metal 
contains  carbon,  but  in  smaller  quantity  than  when  pure  car- 
bonic oxide  is  the  reducing  agent. 

EXP.  109. — Reduce  about  2  grams  of  ferric  oxide  in  the  same 
way  i-hat  lead  oxide  is  reduced  in  Exp.  105,  and  compare  the  result 
of  heating  1  gram  of  the  reduced  metal  in  a  porcelain  crucible  with 
the  similar  results  obtained  as  described  above. 


REDUCTION  193 

The  complex  character  of  coal  gas  makes  the  individual 
reactions  somewhat  numerous  when  it  is  used  as  a  reducing 
agent,  but  the  final  result  is  the  same  as  would  be  obtained 
with  carbon  and  hydrogen  together.  The  reaction  with  the 
principal  hydrocarbon  when  lead  oxide  is  reduced  by  coal  gas 
is  shown  by  the  equation  : 

4PbO  +  CH4  =  4Pb  +  C02  +  2H20. 

Marsh  gas. 

Reduction  by  Metals.— The  great  affinity  which  a  par- 
ticular metal  has  for  an  element,  or  group  of  elements,  in  a 
compound  often  enables  it  to  replace  another  metal  already  in 
the  compound,  and  so  liberate  it  in  the  metallic  state,  when 
the  proper  conditions  are  present.  This  is  the  case  with 
aluminium,  which  exerts  a  powerful  reducing  action  on  some 
metallic  oxides  in  consequence.  The  action  is  accompanied  by 
an  extremely  rapid  evolution  of  heat,  and  the  local  tempera- 
ture obtained  is  sufficient  to  melt  the  reduced  metal,  even  in 
the  case  of  such  difficultly  fusible  metals  as  iron  and  chromium. 
Experiments  illustrating  this  mode  of  reduction  are  easily 
carried  out.  The  aluminium  used  must  be  in  a  finely-divided 
state.  Filings,  which  are  readily  made  from  a  piece  of  the 
rod  metal  with  a  moderately  coarse  file,  are  suitable  for  opera- 
tions on  the  small  scale.  The  oxides  in  the  form  of  fine 
powder  should  be  heated  before  use  to  insure  their  freedom 
from  moisture. 

EXP.  110. — Well  mix  together  10  grams  of  dry  black  oxide  of 
copper  and  3  grams  of  aluminium  filings  ;  put  the  mixture  into  a 
small  clay  crucible ;  make  a  small  cavity  in  the  centre  of  the 
mixture  with  a  rod,  and  into  this  pour  a  mixture  of  2  grams  of 
barium  peroxide  and  O'l  gram  of  aluminium  filings.  Then  insert 
in  this  a  piece  of  magnesium  ribbon  about  2  inches  long  in  an 
upright  position.  Place  the  pot  on  a  brick,  ignite  the  ribbon  with 
a  lighted  taper,  and  rapidly  replace  the  lid.  When  the  reaction  is 
finished  and  the  crucible  has  cooled  down,  break  it  and  extract  the 
button  of  metallic  copper,  which  is  usually  buried  in  a  mass  of  slag 
at  the  bottom  of  the  pot.  Repeat  the  experiment  with  a  mixture  of 
5  grams  of  dry  ferric  oxide  and  2  grams  of  aluminium  filings.  The 

13 


194  METALLURGICAL  CHEMISTRY 

same  quantity  of  the  barium  peroxide  mixture  may  be  used.    When 
the  pot  has  cooled,  break  it  and  extract  the  button  of  metal. 

The  change  is  a  comparatively  simple  one,  for  it  consists  of 
the  oxidation  of  one  metal  at  the  expense  of  the  oxygen  in  the 
oxide  of  the  other. 

3CuO  +  2  Al  =  A1203  +  3Cu. 
237        54 

Fe203  +  2  Al  =  A1203  +  2Fe. 
160       54 

A  slight  excess  of  aluminium  over  the  quantities  found  by 
calculation  from  the  equations  is  used  in  the  experiments 
described  above,  and  part  of  this  excess  passes  into  the  reduced 
metal.  If  in  the  case  of  the  ferric  oxide  an  excess  of  that 
compound  is  used,  the  reduction  is  only  partial,  and  the  tem- 
perature obtained  correspondingly  lower.  The  mixture  may 
be  arranged  for  reduction  to  ferrous  oxide,  thus  : 

3Fe203  +  2  Al  =  6FeO  +  A1203. 

When  this  mixture  is  packed  round  the  prepared  joint  of  a 
steel  tube,  which  is  to  be  hard  soldered,  and  fired,  the  metal 
is  very  rapidly  raised  to  the  temperature  necessary  to  melt 
the  solder  and  make  a  sound  joint. 

This  method  of  extracting  metals  from  their  oxides  is  useful 
in  the  case  of  such  oxides  as  those  of  chromium  and  manganese, 
when  comparatively  small  quantities  of  these  metals  are 
required  for  special  purposes.  The  aluminium  used  for  their 
reduction  must  be  very  finely  divided,  or  the  reduction  is 
incomplete. 

The  reducing  action  of  metallic  sodium  when  brought  into 
contact  with  fused  chlorides  of  metals  is  very  marked,  and 
was  used  for  many  years  in  the  last  stage  of  the  extraction 
of  aluminium  on  the  large  scale.  But  the  method  has  now 
been  superseded  by  the  use  of  the  electric  current  for  the 
extraction  of  the  metal.  The  reaction  depends  upon  the 


REDUCTION  195 

readiness  with  which  sodium  combines  with  chlorine  to  form 
common  salt.     The  equation  is  : 

A1C18  +  3Na  =  Al  +  3NaCl. 

Aluminium 
chloride. 

Potassium  has  similar  reducing  properties. 

Some  of  the  common  metals  will  reduce  others  from  thejr 
combinations  with  sulphur  when  they  are  brought  into  contact 
with  the  fused  sulphides.  The  reaction  with  iron  and  lead 
sulphide  is  a  good  example  of  this  kind  of  change  ;  but  the 
mixture  must  be  raised  to  a  good  red  heat  for  effective 
separation. 


Reference  should  also  be  made  to  Exp.  70,  and  similar  experi- 
ments in  Chap.  VIII.,  as  examples  of  reduction  in  which  the 
metal  is  displaced  from  its  combination  with  a  group  of 
elements.  The  reaction  for  the  separation  of  copper  from 
copper  sulphate  by  zinc  is  expressed  thus  : 

CuS04  +  Zn  =  ZnS04  +  Cu. 

Similar  reactions  are  of  great  importance  in  a  number  of 
metallurgical  operations.  In  such  cases  the  cheapest  suitable 
material  is  employed.  Iron  is  much  used. 

Reduction  by  Potassium  Cyanide.—  Simple  reductions 

can  be  brought  about  by  heating  some  oxides  and  sulphides 
with  potassium  cyanide,  KCN,  at  a  moderate  temperature. 
Great  care  must  be  taken  in  handling  the  cyanide,  on  account 
of  its  highly  poisonous  nature. 

EXP.  111.  —  Put  3  grams  of  potassium  cyanide  into  a  porcelain 
crucible,  and  melt  it  over  a  good  Bunsen  flame.  When  thoroughly 
melted,  add  1  gram  of  tin  oxide,  put  on  the  lid,  and  continue  the 
heating  for  fifteen  minutes.  If  the  globules  of  reduced  metal  have 
not  run  together,  grip  the  side  of  the  open  crucible  with  a  pair 
of  tongs,  and  gently  shake  it  over  the  flame  until  a  single  globule 
of  metal  is  obtained.  Pour  the  liquid  contents  of  the  crucible  on  to 
a  dry  porcelain  slab,  and,  when  solid,  detach  the  excess  of  cyanide 
and  weigh  the  button  of  metal. 

13—2 


196  METALLUKGICAL  CHEMISTRY 

EXAMPLE.—  The  button  of  metal  from  1  gram  of  oxide  weighed 
0-782  gram. 

/.  1  —  0*782=0-218  gram  of  oxygen  removed. 


Then  =  0-0066;  =  0-0136. 

llo  lo 

What  is  the  formula  of  the  tin  oxide  used?  (See  p.  97) 
The  experiment  by  which  this  result  was  obtained  was  made 
with  practically  pure  tin  oxide.  The  ordinary  compound 
(putty  powder)  is  usually  contaminated  with  oxide  of  lead, 
and  gives  too  high  a  result  for  the  metal.  The  reaction  is 
thus  expressed  : 

Sn02  +  2KCN  =  Sn  +  2KCNO. 

Potassium  Potassium 

cyanide.  cyanate. 

Sulphide  of  antimony  (stibnite)  may  be  reduced  in  a  similar 
manner  ;  but  it  is  usually  too  impure  to  give  a  good  quantita- 
tive result.  Oxide  of  lead  is  also  easily  reduced  by  the  molten 
cyanide. 

Reduction  by  Reaction.—  This  kind  of  reduction  is 
largely  used  in  the  separation  of  copper  and  lead  from  their 
ores,  and  is  very  interesting  on  that  account.  It  depends 
upon  the  fact  that  if  a  mixture  of  compounds,  containing  the 
elements  of  another  compound,  which  can  exist  in  the  gaseous 
state  at  a  temperature  below  that  at  which  the  remaining 
elements  in  the  mixture  are  volatilized,  is  heated,  the  gas 
forms  and  escapes  from  the  mixture.  The  non-volatile  element 
is  left  behind  in  the  free  state. 

EXP.  112.  —  Mix  together  1  gram  of  copper  oxide,  CuO,  and  1  gram 
of  copper  sulphide,  Cu2S.  Put  the  mixture  into  a  small  test-tube, 
and  heat  it  strongly  in  the  Bunsen  flame.  Note  that  sulphur 
dioxide  is  given  off,  and  examine  the  residue  when  cold  for  particles 
of  finely  divided  copper. 

The  reaction  is  expressed  by  the  equation  : 


EXP.  113.  —  Mix  together  1  gram  of  lead  oxide  and  1  gram  of  lead 
sulphate,  and  heat  as  above. 


REDUCTION  197 

The  liberated  sulphur  dioxide  is  readily  detected  by  its 
odour,  but  the  reduced  metal  is  not  so  easily  seen  as  the 
copper  in  the  residue  from  the  last  experiment.  The  equa- 
tion is  : 


The  value  of  these  experiments  is  in  the  direct  proof  they 
furnish  of  the  formation  of  sulphur  dioxide.  In  ordinary 
metallurgical  experiments  the  separation  of  the  metals  is 
taken  as  an  indirect  proof  of  the  way  in  which  the  reduction  is 
effected. 

Reduction  Step  by  Step.  —  It  is  possible  to  reduce  some 
compounds  in  stages,  and  thus  prove  a  definite  relation  between 
the  proportions  of  two  elements  in  two  or  more  distinct  com- 
pounds of  the  same. 

EXP  114.  —  Weigh  2  grains  of  lead  peroxide,  Pb02  ;  transfer  it  to 
the  combustion-tube,  Fig.  35,  and  weigh  the  whole.  Heat  the 
powder  until  it  turns  yellow,  but  do  not  let  it  fuse.  Allow  the  tube 
to  cool,  and  reweigh  it.  The  loss  is  due  to  the  escape  of  oxygen 
expelled  by  heat.  Replace  the  tube  in  the  clip,  connect  the  end 
with  the  gas-tap,  and  reduce  the  yellow  powder  to  metallic  lea  din  a 
stream  of  coal-gas.  When  the  reduction  is  finished  and  the  tube  is 
cold  reweigh  it.  The  loss  is  due  to  the  removal  of  oxygen  abstracted 
by  the  constituents  of  the  coal  gas. 

EXAMPLE.  —  In  a  carefully  worked  experiment  2  grams  of  the 
peroxide  lost  0*133  gram  on  heating,  and  0*129  gram  on  reduction 
in  coal-gas.  Therefore  2  —  0'262  =  1*738  grams  of  lead  are  united 
with  0133  and  0*129  gram  of  oxygen  respectively.  If  these  numbers 
are  divided  by  the  atomic  weights  of  the  elements  to  which  they 
belong,  the  ratio  1  :  1  :  1  is  obtained. 

This  proves  that  the  oxygen  is  removed  in  two  equal  steps, 
and  makes  it  evident  that  the  second  compound  may  be 
regarded  as  formed  from  the  first  by  the  addition  of  an  amount 
of  oxygen  strictly  proportional  to  its  atomic  weight.  This  is 
also  true  of  similar  series  of  compounds.  Numerous  examples 
might  be  brought  forward,  and  they  would  all  bear  testimony 
in  the  same  direction.  The  following  are  examples  of  such 
series  :  CO,  C02  ;  Cu20,  CuO  ;  CrO,  Cr203,  Cr03. 


198  METALLURGICAL  CHEMISTRY 

The  generalization  based  upon  these  facts  is  known  as 
the  law  of  chemical  combination  in  multiple  proportion  (see 
p.  52). 

The  series  of  compounds  included  in  this  generalization, 
which,  it  must  be  remembered,  is  based  entirely  upon  experi- 
mental evidence,  and  is  therefore  a  statement  of  facts,  afford 
very  strong  support  to  the  atomic  theory.  For  it  is  perfectly 
clear  that  after  a  particular  compound  has  been  formed,  a 
second  addition  of  one  of  its  elements  may  take  place,  by 
which  another  equally  definite  compound  is  produced. 
Further,  the  amount  of  the  added  element  is  strictly  propor- 
tional to  its  atomic  weight,  for  the  second  compound  is 
formed  by  a  distinct  jump  from  one  proportion  to  the  other. 
It  is,  then,  easy  to  imagine  that  the  molecules  of  the  higher 
compound  are  formed  by  the  addition  of  invariable  particles 
or  atoms  of  the  element  to  the  molecules  of  the  lower 
compound. 

SUMMARY. 

Reduction  is  one  of  the  most  important  operations  in 
metallurgical  work.  The  selling  price  of  a  metal  depends 
largely  upon  the  readiness  with  which  it  is  reduced  from  its 
ores.  The  reduction  of  a  metallic  compound  may  be  complete, 
and  the  metal  set  free ;  or  it  may  be  only  partial,  and  a 
lower  compound  of  the  metal  obtained.  Only  a  few  metals 
are  liberated  from  their  compounds  by  heat  alone.  Assistance 
has  to  be  rendered  by  reducing  agents,  of  which  carbon, 
carbonic  oxide,  and  hydrogen  are  the  most  important  for  the 
reduction  of  metallic  oxides.  The  formation  of  carbon 
dioxide  and  water  is  easily  proved.  The  electric  current 
is  coming  more  and  more  into  prominence  as  a  reducing 
agent,  and  it  is  in  use  on  the  large  scale  for  the  extraction 
of  several  metals.  Metals  themselves  are  not  quite  so 
prominent  in  the  list  of  reducing  agents  as  they  were,  but 
iron,  copper,  and  zinc  are  still  largely  used  in  the  metallurgy 
of  gold  and  silver.  Potassium  cyanide  is  only  used  on  the 


KEDUCTION  199 

small  scale  for  the  reduction  of  metals  from  oxides  and 
sulphides.  Step-by-step  reduction,  and  its  bearing  on  the 
second  law  of  chemical  combination,  is  of  great  theoretical 
importance. 

QUESTIONS. 

1.  What  do  you  understand  by  the  partial  and  complete 
reduction  of  a  compound  ?     Give  examples. 

2.  How  may  a  reduction  experiment  be  used  to  determine 
the  formula  of  an  oxide  1 

3.  Describe  an  experiment  by  which  the    composition  of 
water  is  determined. 

4.  What  is  the  composition  of  coal  gas  T    Write  out  equa- 
tions showing  the  effect  of  its  various  constituents  upon  a 
heated  metallic  oxide. 

5.  Write  down  the  law  of  chemical  combination  in  multiple 
proportions.     What  is  the  nature  of  the  support  it  lends  to 
the  atomic  theory  ? 

6.  How  is  sulphur  dioxide  formed  in  a  reduction  process 


CHAPTER  XII 
COMBUSTION 

IN  Chap.  XL  some  experiments  are  described  in  which  metals 
are  reduced  from  their  oxides  by  carbon  and  hydrogen  re- 
spectively ;  and  it  is  proved  that  carbon  dioxide  and  water 
are  formed  in  the  reactions  which  take  place.  From  this  it 
is  evident  that  metallic  oxides  can  act  as  oxidizing  agents,  and 
in  doing  so  give  up  their  oxygen  to  the  bodies  by  which  they 
are  reduced  to  the  metallic  state.  But  in  order  that  these 
changes  may  take  place,  it  is  necessary  to  heat  the  metallic 
oxides  in  contact  with  the  reducing  agents  above  certain 
limiting  temperatures,  below  which  no  action  is  observed. 
The  reason  for  this  is  that  the  affinity  of  the  oxygen  for  the 
metal  must  be  overcome  before  the  carbon  or  hydrogen  can 
combine  with  it.  The  effect  of  the  decomposing  action  of  the 
heat  on  the  oxide,  and  the  pull  of  the  reducing  agent  upon 
its  oxygen,  bring  about  the  decomposition,  when  they  are 
together  greater  than  the  attraction  of  the  elements  in  the 
oxide  for  each  other.  Then  the  reducing  agents  burn  in  the 
oxygen  of  the  oxides. 

When  carbon  and  hydrogen  oxidize  freely,  and  the  action 
continues  without  the  aid  of  external  heat,  the  change  is 
commonly  known  as  combustion  :  but  the  compounds  formed 
are  the  same  as  in  the  reduction  of  oxides.  The  heat  developed 
in  their  formation  is,  however,  much  greater  than  that  used 
up  in  rendering  the  changes  possible,  and  so  the  action 
continues  as  long  as  the  necessary  materials  are  supplied. 

There  is  a  certain  limiting  temperature  for  a  given  com- 

200 


COMBUSTION  201 

bustible  body,  below  which  it  may  be  exposed  to  the  air 
without  taking  fire.  This  temperature  is  called  the  ignition 
point,  and  varies  for  different  bodies.  If  the  body  is  exposed 
to  the  air  at  this  temperature,  or  a  little  above  it,  oxidation 
commences,  heat  is  developed,  and  a  further  rise  in  tempera- 
ture takes  place.  This  increases  the  action,  and  if  the  body 
burns  freely  its  maximum  rate  of  combustion  is  soon  reached, 
and  continues  steadily  as  long  as  the  conditions  are  constant. 
The  rate  of  combustion  must  depend  upon  the  rate  at  which 
the  particles  of  oxygen  are  brought  into  contact  with  the 
combustible  body.  Also,  the  local  temperature  depends 
largely  upon  the  rate  of  combustion,  and  this  can  be  increased 
by  bringing  more  particles  of  oxygen  into  contact  with  the 
burning  body  in  a  given  time.  This  is  commonly  effected  by 
blowing,  and  the  higher  temperature  obtained  by  burning  the 
combustible  body  in  a  rapid  current  of  air  is  explained.  A 
further  increase  in  the  temperature  takes  place  if  the  air  is 
heated  before  it  comes  into  contact  with  the  burning  body  ; 
for  the  heat  which  would  be  used  up  in  raising  the  products 
of  combustion  and  the  residual  nitrogen  to  the  temperature 
of  the  heated  air,  is  now  available  for  increasing  the  general 
temperature. 

It  has  been  already  stated  that  the  quantity  of  heat  liberated 
during  a  chemical  combination  is  just  as  definite  as  the 
quantity  of  the  compound  formed.  Thus  the  thermo-chemical 
equation 


12    32      44 

means  that  12  grams  of  carbon  in  burning  completely  to 
carbon  dioxide  liberates  sufficient  heat  to  raise  97  kilograms 
of  water  through  1°  C.  Similar  thermal  equations  may  be 
written  for  other  combustible  bodies. 

Further,  it  must  be  carefully  noted  that  exactly  the  same 
quantity  of  heat  disappears  if  the  compound  formed  during 
combustion  is  decomposed  into  its  elements  again.  And  it  is 
also  certain  that  there  is  a  limiting  temperature  at  which  the 


202  METALLUKGICAL  CHEMISTKY 

compound  is  decomposed  as  fast  as  it  is  formed.  This  higher 
limit  is  the  temperature  of  dissociation  of  the  compound,  and 
is  the  theoretical  limit  to  the  temperature  obtained  by  the 
combustion  of  the  element  forming  the  compound.  The 
practical  limit,  which  is  influenced  by  a  variety  of  circum- 
stances, seldom  reaches  this.  Some  assistance  in  understand- 
ing this  limit  may  be  obtained  by  considering  the  experiment 
with  mercury,  described  on  p.  17.  Here  it  is  evident  that  a 
certain  minimum  temperature  is  necessary  in  order  that  the 
mercury  may  oxidize ;  but,  if  the  oxide  is  exposed  to  a  tem- 
perature not  much  higher  than  that  at  which  it  is  formed, 
decomposition  sets  in,  and  the  elements  are  reproduced.  It  is, 
then,  certain  that  there  is  a  temperature  at  which  the  oxide  is 
decomposed  as  fast  as  it  is  formed,  and  that  above  this  limit 
no  permanent  oxidation  is  possible.  Thus,  however  favourable 
the  conditions  may  be  under  which  a  body  is  burnt,  there  is  a 
limit  to  the  temperature  of  the  products  of  its  combustion. 
When  air  is  used  for  burning  a  body,  every  molecule  of 
oxygen  brought  into  contact  with  it  has  four  molecules  of 
nitrogen  in  its  train.  Now,  if  these  are  replaced  by  oxygen 
molecules,  it  is  evident  that  a  more  rapid  action  must  result. 
This  is  what  happens  when  pure  oxygen  is  used  in  place  of 
air  :  it  has  the  same  general  effect  as  a  rapid  current  of  air, 
without  the  disadvantage  of  the  presence  of  a  considerable 
body  of  nitrogen  to  absorb  heat,  and  thus  to  lower  the  general 
temperature.  If,  then,  pure  oxygen  is  used  to  surround  the 
burning  body  instead  of  air,  the  effect  should  be  considerably 
increased.  This  may  be  rendered  perfectly  clear  by  making  a 
few  experiments  with  oxygen  itself. 

OXYGEN. 

Preparation  of  Oxygen  from  Air  (Erin's  process).— 
This  commercial  process  depends  upon  the  fact  that  when 
barium  oxide,  BaO,  is  heated  to  a  certain  temperature  in  a 
current  of  air,  it  absorbs  oxygen  rapidly,  and  is  converted  into 


COMBUSTION  203 

* 

the  peroxide,  Ba02,  and  that  when  this  compound  is  heated 
to  a  still  higher  temperature,  it  is  decomposed  into  the  lower 
oxide  again,  oxygen  being  set  free.  It  is  another  case  of  a 
limiting  temperature,  above  which  the  change  is  reversed. 
But  it  is  found  that  by  reducing  the  pressure  over  the  hot 
peroxide  the  limiting  temperature  for  decomposition  may  be 
brought  down  to  that  at  which  absorption  of  oxygen  takes 
place  under  ordinary  pressure.  By  taking  advantage  of  this 
the  process  is  made  an  alternating  one,  by  varying  the 
pressure  instead  of  the  temperature,  which  is  much  more 
convenient. 

Barium  oxide  is  heated  in  iron  tubes  arranged  in  a  furnace, 
and  a  current  of  air,  freed  from  carbon  dioxide  and  water  by 
passing  it  through  vessels  containing  caustic  soda  and  lime, 
is  passed  over  it.  Absorption  takes  place  : 


Barium  Barium 

oxide.  peroxide. 

The  air  current  is  then  stopped,  and  the  tubes  carefully 
exhausted  by  a  pump,  when,  without  interfering  with  the 
temperature  of  the  tubes,  oxygen  gas  is  liberated. 

2Ba02  =  2BaO  +  02. 

The  first  portion  of  the  gas  is  allowed  to  escape,  as  it  is  con- 
taminated with  the  nitrogen  present  in  the  tubes  when  the 
air  is  cut  off  ;  the  remainder,  which  is  pure  oxygen,  is  passed 
into  a  gas-holder  as  fast  as  it  is  liberated.  The  oxygen  thus 
obtained  is  forced  in  large  quantities  into  strong  portable  iron 
bottles,  from  which  it  can  be  drawn  at  will.  The  gas  stored 
in  this  way  is  very  convenient  for  use  at  short  intervals. 

Preparation  of  Oxygen  from  Potassium  Chlorate. 

—  When  small  quantities  of  the  gas  are  required  at  irregular 
intervals,  it  is  better  to  prepare  it  as  required. 

EXP.  115.  —  Mix  10  to  20  grams  of  powdered  potassium  chlorate  with 
a  quarter  its  weight  of  black  oxide  of  manganese,  and  put  the  mixture 
into  a  test-tube,  1  inch  wide  and  7  inches  long,  fitted  with  a  bung  and 


204  METALLUKGICAL  CHEMISTKY 

delivery-tube.  Arrange  the  tube  in  a  horizontal  position,  as  shown  on 
p.  39,  so  that  there  is  a  clear  channel  between  the  mixture  and  the  side 
of  the  tube,  and  with  the  free  end  of  the  mixture  not  too  near  the 
bung.  Heat  the  tube  carefully,  beginning  at  the  free  end  of  the 
mixture,  and  gradually  carry  the  flame  along  as  the  escape  of  gas 
slows  down.  Collect  several  jars  of  the  gas,  and  one  stoppered  bell 
jar  of  about  1,500  c.c.  capacity. 

One  gram  of  the  chlorate  liberates  about  270  c.c.  of  oxygen, 
so  that  a  rough  calculation  can  be  made  for  the  quantity 
required,  and  waste  thus  avoided.  The  change  is  given  by 
the  equation : 

2KC103   =    2KC1  +  302. 

Potassium  Potassium 

chlorate.  chloride. 

The  black  oxide  undergoes  very  little  change  itself,  but  assists 
the  oxygen  to  escape  at  a  lower  temperature  than  when  the 
chlorate  is  used  alone. 

Properties  Of  Oxygen. — It  is  a  colourless,  odourless,  and 
tasteless  gas ;  only  slightly  soluble  in  water,  and  the  most 
important  supporter  of  combustion.  The  following  experi- 
ments may  be  made  with  the  gas  prepared  as  above  : 

1.  When  a  glowing  splint  is  put  into  the  gas  it  is  rekindled. 

2.  When  a  piece  of  good  charcoal  is  made  red  hot  in  the 
Bunsen  flame,  and  then  plunged  into  the  gas,  it  burns  vigor- 
ously, emitting  much  heat  and  light,  but  without  flame.     The 
residual  gas  turns  lime-water  milky ;  it  is  carbon  dioxide. 

3.  Sulphur  when  feebly  ignited  bursts  into  more  vigorous 
combustion,  and  burns  with  a  violet  flame.     The  residual  gas 
has  the  odour  of  burning  sulphur,  and  turns  litmus  red.     It 
is  the  acid-forming  sulphur  dioxide. 

4.  When  ignited  phosphorus  is  put  into  oxygen  it  burns 
with  a  very  bright  flame.     The  acid  oxide  P205  is  formed. 
A  common  jar  should  be  used  for  this  experiment,  as  it  is  liable 
to  fracture  from  the  great  heat  developed. 

5.  A  piece  of  thin  steel  ribbon  made  into  a  spiral,  and  one 
end  tipped  with  a  short  piece  of  a  wax  match,  burns  very 
vividly  when  the  wax  is  ignited  and  the  whole  put  into  a 


COMBUSTION 


205 


bell -jar  filled  with  oxygen,  and  standing  over  water.  See 
Fig.  38. 

6.  If  two  small  candles  supported  in  deflagrating  spoons  are 
ignited  and  put  into  two  cylinders, 
one  containing  air  and  the  other 
oxygen,  a  very  marked  difference 
is  observed  in  both  the  rate  of 
combustion  and  the  appearance  of 
the  flames.  In  the  air  the  candle 
burns  at  the  normal  rate  for  a 
short  time,  then  the  flame  lan- 
guishes, and  finally  disappears.  In 
the  oxygen  the  flame  shrinks,  be- 
comes much  brighter,  and  the  rate 
of  combustion  is  increased.  An 

increase  in  temperature  is  evident  from  the  fact  that  the  wax 
melts,  and  runs  down  into  the  spoon.  The  flame  finally  dis- 
appears. The  formation  of  carbon  dioxide  during  the  com- 
bustion is  easily  proved  by  pouring  a  little  clear  lime-water 

into  each  jar;  on  shaking,  the  lime-water  is  turned 

milky. 

Combustion  in  Air  and  other  Bodies.— The 

burning  of  a  candle  may  be  made  the  subject  of  an 
experiment  by  which  the  increase  in  weight  due  to 
the  combination  of  oxygen  with  the  constituents  of 
the   combustible   body   may   be    proved.      A   lamp 
chimney,  Fig.  39,  is  divided  into  two  parts  by  a  piece 
of  wire  gauze.     The  upper  part  is  filled  with  a  mix- 
ture of  caustic  soda  and  calcium  chloride  in  small 
lumps,  arid  the  upper  end  is  fitted  with  a  perforated 
cork.    Another  perforated  cork  bearing  a  small  candle 
is  inserted  in  the  lower  end.    The  whole  is  then  hung 
by  a  wire  loop  on  to  the  end  of  the  scale-beam  of  a  balance 
from  which   the  scale-pan  has   been   removed,   and  counter- 
poised.    The  lower  cork  is  then  removed,  the  candle  ignited, 


206  METALLURGICAL  CHEMISTRY 

and  rapidly  returned  to  the  tube.  Very  soon  the  apparatus 
begins  to  show  an  increase  in  weight,  although  the  candle  is 
actually  disappearing.  The  cause  of  this  increase  is  readily 
explained,  if  it  is  borne  in  mind  that  the  carbon  and  hydrogen 
of  the  candle  combine  with  and  thus  abstract  oxygen  from  the 
air  passing  through  the  apparatus,  and  that  the  absorbing 
material  in  the  upper  part  prevents  the  escape  of  the  pro- 
ducts of  the  combustion.  Thus  the  whole  of  the  matter  of  the 
candle  is  kept  in  the  tube,  together  with  the  added  oxygen. 
This  is  evidently  the  cause  of  the  increase  in  weight. 

The  formation  of  water  during  the  combustion  of  hydrogen, 
or  combustible  bodies  containing  hydrogen,  is  easily  demon- 
strated. The  drying-tube  of  the  part  of  the  apparatus  for 
generating  dry  hydrogen  is  fitted  with  a  piece  of  clay  tobacco- 
pipe,  and  the  long  tube  passing  into  the  wash-bottle  is  raised 
just  above  the  surface  of  the  liquid  (p.  189).  The  apparatus 
shown  in  Fig.  19  is  carefully  dried,  and  the  head  of  the  thistle 
funnel  arranged  over  the  clay  jet,  the  angle-tube  being  con- 
nected with  an  aspirator.  The  jet  of  hydrogen  is  then  ignited, 
and  a  current  of  air  aspirated  through  the  thistle  funnel. 
Very  soon  water  appears  in  the  bottle,  and  the  only  explana- 
tion of  its  presence  there  is  that  it  has  been  formed  by  the 
burning  of  the  hydrogen  in  air.  A  jet  of  coal  gas  or  a  burning 
candle  may  be  placed  under  the  funnel  with  a  similar  result. 

Combustion  is  not  limited  to  air  or  oxygen,  for  hydrogen^ 
combustible  bodies  containing  hydrogen,  and  finely  divided 
metals  burn  readily  in  chlorine  (p.  74).  Nitrous  oxide  and 
nitric  peroxide  are  also  supporters  of  combustion.  Finely 
divided  metals  will  also  burn  in  sulphur  vapour. 

EXP.  116.—  Put  some  sulphur  into  a  wide  test-tube  ;  fix  the  tube 
in  an  upright  position,  and  strongly  heat  it.  When  the  sulphur  is 
boiling  vigorously  lower  a  spiral  of  thin  copper  ribbon  into  the 
vapour.  In  a  short  time  it  becomes  red  hot,  due  to  the  rapid 
sulphidation.  The  metal  burns  in  the  sulphur  vapour.  When  the 
spiral  is  withdrawn  it  is  fouod  to  be  changed  into  copper  sulphide. 

Explosive    Combustion.— In    ordinary   combustion   the 


COMBUSTION  207 

conditions  are  such  that  the  burning  body  and  the  supporter 
of  combustion  are  supplied  to  each  other  steadily,  and  any 
alteration  in  the  rate  of  supply  alters  the  rate  at  which  the 
action  takes  place.  Explosion  of  a  mixture  of  gases  or  solids 
is  simply  rapid  combustion,  and  takes  place  when  the  two 
bodies  are  so  intimately  mixed  that  the  entire  action  is  spread 
over  a  very  short  space  of  time. 

The  explosion  of  a  mixture  of  coal  gas  and  air  is  a  familiar 
example,  and  the  energy  of  the  change  is  unmistakable. 
The  working  of  an  ordinary  gas  engine  depends  upon  a  series 
of  such  explosions.  Gunpowder  is  an  example  of  an  explosive 
mixture  of  solids  ;  it  is  a  case  of  the  rapid  oxidation  of  carbon 
and  sulphur  by  the  oxygen  obtained  from  the  saltpetre.  A 
large  quantity  of  gaseous  matter  is  thus  formed  in  a  limited 
space,  and  the  enormous  pressure  which  this  gas 
is  able  to  exert  easily  accounts  for  the  effects 
produced. 


Inversion   of  Combustion.  —  The   terms 

combustible  and  supporter  of  combustion  are 
merely  relative.  In  the  ordinary  way  we  say 
that  coal  gas  burns  in  air ;  but  it  is  easy  to 
invert  this  order  of  things,  and  make  the  air 
burn  in  coal  gas.  A  very  simple  experiment 
will  prove  this. 

EXP.  117.— Fix  a  lamp  chimney  in  a  clip,  insert  a 
cork,  through  which  two  tubes  pass,  as  shown  in 
Fig.  40,  in  the  bottom,  and  put  a  gauze  cap  on  the 
top.  Connect  the  angle-tube  a  with  the  gas-tap,  and 
ignite  the  jet  of  gas  escaping  through  the  gauze.  Tip 
the  end  of  a  thin  splint  of  wood  with  a  fragment  of 
phosphorus,  ignite  it,  pass  the  splint  rapidly  up  the 
straight  tube  b,  and  ignite  the  jet  of  air  at  the  top.  FlG-  40- 

The  direct  ignition  of  the  jet  of  air  as  described,  although 
more  convincing,  is  not  always  easy  to  bring  about.  Another 
way  is  to  close  the  top  of  the  chimney  with  a  cork,  through 
which  a  glass  tube  is  passed,  to  turn  on  the  gas  freely,  and 


208  METALLUEGICAL  CHEMISTEY 

to  ignite  it  at  both  top  and  bottom  tubes.  If  then  the 
supply  is  regulated  carefully,  the  flame  at  the  end  of  the 
bottom  tube  will  strike  back  up  the  tube,  and  appear  at  the 
inner  end  as  a  jet  of  air  burning  in  coal  gas. 

The  current  of  coal  gas  passing  in  through  a  and  up  the 
chimney  draws  in  a  current  of  air  through  b,  so  that  a  jet  of 
air  is  supplied  to  an  atmosphere  of  coal  gas.  The  air  takes 
fire  and  burns  with  a  flame,  which,  though  not  so  luminous  as 
the  flame  of  coal  gas  burning  in  air,  resembles  it  in  general 
appearance. 

Similarly,  oxygen  may  be  burnt  in  hydrogen,  and  chlorine 
in  hydrogen.  It  can  then  be  said  that  when  two  gases  are 
concerned  in  the  act  of  burning  either  may  be  supplied  in  a 
jet  to  the  other  as  a  supporting  atmosphere.  But  the  chemical 
change  is  the  same  whichever  way  the  action  takes  place. 
Thus,  for  oxygen  and  hydrogen  water  is  the  invariable  pro- 
duct. If,  however,  one  of  the  bodies  is  a  non-volatile  solid, 
the  inversion  is  not  possible. 

Flame. — The  burning  of  carbon  and  sulphur  in  oxygen 
furnishes  examples  of  combustion  in  which  a  marked  difference 
is  noticed.  The  carbon  becomes  incandescent,  and  gives  out 
light  and  heat,  but  without  flame.  It  is  a  non-volatile  solid. 
The  sulphur,  on  the  other  hand,  gets  hot  and  vaporizes.  This 
vapour  diffuses,  penetrates  the  surrounding  space  for  some 
distance,  and  burns  there.  The  space  in  which  the  vapour 
burns  is  filled  with  flame,  and  in  the  whole  of  this  region  the 
combustion  is  going  on.  In  the  first  case  the  zone  of  combustion 
is  confined  to  the  surface  of  the  burning  solid,  and  the  heat 
developed  is  concentrated  in  space.  In  the  second  the  zone 
of  combustion  is  extended  to  the  region  about  the  burning 
body,  and  the  heat  is  distributed  over  a  larger  space. 

The  appearance  of  the  space  in  which  gaseous  matter  is 
burning  gives  the  general  idea  of  flame,  and  the  further  the 
vapour  penetrates  into  the  space  before  it  is  completely  burnt 
the  larger  will  be  the  volume  of  the  flame.  Bearing  this  in 


COMBUSTION  209 

mind  it  is  easy  to  explain  why  a  candle  flame  shrinks  when  it 
is  fed  with  pure  oxygen.  For  the  oxygen  burns  up  the  com- 
bustible gases  emanating  from  the  wick  of  the  candle  before 
they  have  time  to  penetrate  very  far  into  the  surrounding 
space,  and  a  shrinking  of  the  zone  of  combustion  takes  place. 
The  local  temperature  is  increased  by  this,  and  also  by  the 
more  rapid  burning.  If  it  is  required  to  localize  the  heat  so 
as  to  attain  a  high  temperature  in  a  small  space,  a  combustible 
body  which  preserves  its  solid  state  during  the  whole  period 
of  its  combustion  should  be  used,  and  the  gas  in  which  it  is 
burnt  should  be  supplied  as  fast  as  possible.  On  the  other 
hand,  when  it  is  necessary  to  spread  out  the  heat,  so  to  speak, 
through  a  larger  space,  gaseous  fuel  must  be  used.  Also,  the 
supply  of  air  must  be  such  as  to  allow  of  perfect  combustion 
in  as  large  a  region  as  possible.  Charcoal,  coke,  and  anthra- 
cite are  examples  of  fuel  which,  when  burnt  in  a  good  supply 
of  air,  give  high  local  temperatures.  But  gas  coal,  or  long 
flaming  coal,  which  contains  a  large  percentage  of  volatile 
combustible  matter,  is  more  suitable  for  the  distribution  of 
heat.  Gaseous  fuel,  however,  may  be  used  for  obtaining  a 
high  local  temperature,  if  it  is  suitably  burnt.  The  shrinking 
of  a  candle  flame  in  oxygen  suggests  the  means  to  be  adopted  ; 
and  the  oxy-hydrogen  blow-pipe  flame  is  an  excellent  example 
of  its  application.  The  tendency  to  do  away  with  the  use  of 
solid  fuel  for  all  purposes  to  which  combustible  gas  can  be 
applied  is  growing;  and  the  use  of  gaseous  fuel  may  be 
regarded  as  a  distinct  scientific,  economic,  and  hygienic 
advance  upon  the  dirty,  wasteful  methods  of  heat  production 
still  so  largely  used.  Producer  gas,  water  gas,  and  Mond  gas 
are  now  quite  familiar  terms. 

The  Structure  of  Flame. — The  complexity  of  a  given 
flame  varies  with  the  constituents  of  the  gas  producing  it,  and 
the  conditions  under  which  the  gas  is  burnt.  A  general  idea, 
however,  can  be  obtained  by  a  few  simple  observations.  If 
the  air- way  of  a  Bunsen  burner  is  stopped  a  luminous  flame  is 

14 


210  METALLUEGICAL  CHEMISTRY 

obtained  at  the  top ;  and  if  the  gas  is  turned  down  so  as  to 
give  a  small  flame  without  flickering,  the  flame  is  seen  to 
consist  of  three  parts  :  (1)  a  blue  portion,  appearing  to  spring 
from  the  top  of  the  burner;  (2)  a  luminous  portion,  which 
occupies  the  upper  part  of  the  body  of  the  flame  ;  (3)  a  faintly- 
luminous  mantle,  which  surrounds  the  whole  body  of  the  flame 
and  forms  its  limits.  A  dark  portion,  which  does  not  belong 
to  the  flame  proper,  is  noticed  in  the  lower  part,  and  this 
region  is  filled  with  unburnt  gas.  This  is  easily  proved  by 
thrusting  the  head  of  a  match  rapidly  through  the  envelope 
into  this  space,  and  the  fact  that  it  does  not  ignite  there  is 
sufficient  evidence  that  combustion  is  not  taking  place  in  that 
space.  These  observations  are  most  successful  in  an  other- 
wise dark  room.  An  ordinary  candle  flame  has  an  exactly 
similar  appearance.  The  same  conical  space  filled  with  un- 
burnt gas  is  present  in  the  ordinary  Bunsen  flame,  and  its 
presence  may  be  demonstrated  by  holding  a  piece  of  stiff" 
white  paper  well  down  on  the  flame  for  a  second  or  two,  so 
as  to  cut  through  this  space.  A  black  ring  surrounding  a 
white  centre  shows  where  the  paper  has  been  in  contact  with 
burning  and  non-burning  gas. 

The  Luminosity  of  flame  is  of  considerable  importance 
when  it  is  a  question  of  obtaining  as  much  light  as  possible 
from  the  burning  gas.  The  white  flame,  which  is  so  much 
prized  in  illuminating  gas,  is  not  due  to  any  single  cause. 
It  is  probably  largely  caused  by  the  presence  in  that  portion 
of  the  flame  of  dense  gases  and  solid  particles,  produced  by 
complex  changes  in  the  gas  before  its  complete  combustion. 
An  increase  in  the  pressure  under  which  a  gas  is  burnt  also 
increases  the  luminosity  of  its  flame,  but  this  condition  is  not 
commonly  present  to  any  extent.  The  hydrogen  flame  gives 
out  very  little  light,  which  is  due  to  the  extremely  simple 
character  of  its  combustion;  but  if  fine  particles  of  solid 
matter  are  passed  into  it  its  luminosity  is  at  once  increased. 
This  principle  is  adopted  in  converting  the  practically  lightless 


COMBUSTION  211 

flame  of  a  Bunsen  burner  into  a  brilliant  one  by  suspending 
in  it  a  mantle  made  of  a  very  light  infusible  oxide,  which  is 
rendered  incandescent  by  the  heat  of  the  flame.  The  burning 
of  phosphorus,  on  the  other  hand,  proves  that  it  is  possible  to 
have  a  highly  luminous  flame  without  the  presence  of  solid 
particles  in  it.  For  solid  matter  cannot  possibly  be  present 
in  this  flame.  Also,  the  highly-luminous  flame  of  burning 
acetylene,  C2H2,  is  probably  largely  due  to  the  endothermic 
character  of  the  compound  itself.  Much  heat  is  absorbed  in 
its  formation,  and  this  heat,  which  is  liberated  during  its 
decomposition,  is  added  to  that  evolved  by  the  combustion  of 
its  constituents.  A  very  high  local  temperature  is  thus 
obtained  with  the  other  conditions  for  a  luminous  flame. 

THE  ENERGY  OF  COMBUSTION. 

The  Conservation  of  Energy.— Since  combustion  is 
the  principal  source  of  artificial  heat,  if  it  may  be  so  called 
to  distinguish  it  from  the  natural  heat  derived  from  the  sun, 
it  will  be  in  place  here  to  inquire  a  little  into  the  nature  of 
heat.  In  the  first  place,  heat  is  real,  but  it  is  not  ponderable ; 
that  is,  it  lacks  the  intrinsic  property  of  matter,  which  is  best 
denoted  by  substance  or  density.  It  has  been  defined  as  a 
mode  of  motion  ;  but  as  we  can  form  no  idea  of  motion  with- 
out something  to  move,  it  does  not  appeal  to  us  until  it  is 
associated  with  bodies  of  some  kind.  Heat  motion  in  bodies 
is  a  movement  of  their  particles,  and  not  of  the  bodies  as  a 
whole.  A  full  discussion  of  the  nature  of  heat  will  be  found 
in  books  on  the  subject ;  but,  for  our  present  purpose,  it  is 
sufficient  to  recognise  its  ability  to  cause  an  increase  in  mole- 
cular motion  when  it  enters  bodies,  and  a  decrease  in  that 
motion  when  it  leaves  them. 

When  a  body  is  moving  it  is  said  to  possess  energy, 
for  it  is  capable  of  doing  work.  This  kind  of  energy  is 
kinetic  or  "  moving  "  energy.  Therefore,  since  we  have  no 
experience  of  absolutely  cold  bodies,  such  bodies  as  we  know 

14—2 


212  METALLUKGICAL  CHEMISTKY 

possess  a  store  of  this  moving  energy,  due  to  their  molecular 
motion. 

But  bodies  may  also  possess  energy  due  to  their  position 
with  respect  to  other  bodies.  Thus,  they  may  be  able  to  fall 
and,  in  falling,  to  do  work.  This  kind  of  energy  is  potential, 
or  "  stored  "  energy.  The  potential  energy  of  a  body  appeals 
very  strongly  to  the  mind  at  times.  If  one  is  looking  up  at 
a  piece  of  overhanging  rock  on  the  top  of  a  cliff,  the  feeling 
will  arise'  as  to  what  would  happen  if  it  were  to  fall.  The 
same  piece  of  rock  lying  at  one's  feet  does  not  create  the  same 
impression.  When  a  body  is  falling  its  potential  energy,  which 
depends  upon  its  position,  is  being  rapidly  converted  into 
kinetic  energy,  and  just  before  it  reaches  the  surface  towards 
which  it  is  falling,  all  the  potential  energy  due  to  its  original 
position  with  respect  to  that  surface  is  converted  into  moving 
energy.  By  impact  with  the  surface  the  body  is  brought 
to  rest,  and  it  would  appear  as  if  its  energy  of  motion  were 
destroyed ;  also,  since  it  no  longer  possesses  the  potential 
energy  due  to  its  former  position,  that  a  complete  disappear- 
ance of  energy  had  taken  place.  But  further  examination 
proves  that  the  energy  of  motion  of  the  mass  is  not  destroyed ; 
it  is  converted  into  increased  motion  of  the  particles  of  the 
body  and  the  surface  struck  by  it.  That  is,  it  is  converted 
into  heat.  Further,  the  heat  energy  developed  is  exactly 
equivalent  to  the  kinetic  energy  lost.  The  same  reasoning 
holds  for  every  modification  of  energy.  This  principle,  which 
is  known  as  the  conservation  of  energy,  has  been  most 
exhaustively  studied  in  connection  with  the  relation  between 
heat  and  mechanical  work,  and  in  the  case  of  these  two  modifi- 
cations of  energy  has  been  proved  up  to  the  hilt.  Those  who 
are  most  qualified  to  give  an  opinion  hold  that  the  same  is  true 
for  every  other  modification  of  energy. 

Every  particle  of  matter  has  energy  associated  with  it,  and 
when  a  redistribution  of  the  particles  of  a  body,  or  system  of 
bodies,  takes  place,  there  is  also  a  redistribution  of  the  energy 
associated  with  them.  The  energy  of  chemical  attraction  is  of 


COMBUSTION  213 

the  potential  kind,  and  when  the  conditions  are  such  that  this 
attraction  is  able  to  bring  about  chemical  change  in  the  matter 
of  the  reacting  bodies,  a  change  in  energy  also  usually  results. 
Take,  for  example,  the  case  of  a  mixture  of  oxygen  and 
hydrogen  :  the  two  gases  are  ready  to  combine  with  each  other, 
and  a  considerable  amount  of  energy  is  stored  up  in  their 
molecules.  When  the  mixture  is  exploded  a  redistribution  of 
the  atoms  of  the  gases  to  form  molecules  of  water  takes  place, 
and  this  is  accompanied  by  rapid  conversion  of  the  potential 
energy  of  chemical  attraction  into  the  kinetic  energy  of  the 
heat  evolved.  The  cause  of  the  force  exerted  between 
dissimilar  bodies  when  chemical  changes  are  affected,  and  the 
nature  of  chemical  energy,  are  quite  unknown,  but  this  does 
not  prevent  us  from  using  ideas  suggested  by  the  results.  The 
conversion  of  chemical  energy  into  heat  resembles  very  much 
the  conversion  of  the  energy  of  a  falling  body  into  heat  when 
it  is  brought  to  rest  by  collision  with  the  earth ;  and  it  does 
not  appear  at  all  improbable  that  the  change  in  the  motion  of 
the  atoms  of  oxygen  and  hydrogen^  in  their  endeavour  to  form 
new  and  more  stable  molecules  is  the  cause  of  the  conversion  of 
the  stored  chemical  energy  into  heat.  The  action  going  on 
when  a  piece  of  red-hot  charcoal  is  surrounded  by  oxygen  or 
air  may  be  pictured  in  the  same  way.  The  atoms  of  oxygen 
become  associated  with  the  atoms  of  carbon,  and  form  molecules 
of  carbon  dioxide,  and  the  transformation  of  energy  goes  on 
vigorously.  During  the  rush  and  turmoil  of  the  action, 
chemical  energy  disappears,  and  heat  energy  appears  in  the 
same  proportion. 

No  difficulty  should  now  be  experienced  in  explaining  the 
fact  that  different  modifications  of  carbon  appear  to  give  out 
different  quantities  of  heat  when  they  are  completely  burnt. 
The  different  forms  of  carbon  are  admittedly  due  to  a  different 
arrangement  of  the  molecules  or  atoms  of  the  element,  and  not 
to  a  difference  in  the  quality  of  the  atoms  themselves.  It  is, 
therefore,  quite  feasible  that  the  atoms  of  carbon  in  graphite 
require  more  to  bring  them  up  to  the  condition  necessary  for 


214  METALLURGICAL  CHEMISTEY 

union  with  oxygen  than  do  the  similar  atoms  of  carbon  in 
charcoal,  and  that  the  energy  required  to  do  this  work  of 
preparation  is  subtracted  from  the  total  quantity  of  heat, 
which  would  be  developed  by  the  burning  of  the  free  atoms  of 
carbon.  In  fact,  it  is  analogous  to  the  case  of  a  body  which 
has  to  be  lifted  into  a  position  from  which  it  can  fall.  Here  the 
energy  used  up  in  lifting  the  body  into  position  is  to  be 
subtracted  from  the  total  heat  developed  by  the  impact  of  the 
body  when  it  is  brought  to  rest.  It  may,  then,  be  said  that  the 
distribution  of  energy  in  the  allotropic  modifications  of  carbon 
is  not  the  same.  This  holds  for  sulphur  too  ;  for  the  monoclinic 
form  of  that  element  evolves  heat  when  its  crystals  pass  into 
the  more  stable  rhombohedrons  (p.  57).  Generally,  when  an 
element  passes  from  a  more  stable  to  a  less  stable  form,  heat  is 
absorbed,  but  heat  is  evolved  in  the  opposite  change.  When  a 
group  of  molecules  of  different  bodies  are  reacting  among 
themselves,  and  forming  molecules  of  a  new  body  or  bodies, 
evolution  or  absorption  of  heat  takes  place  according  as  the 
new  molecules  are  more  or  less  stable  than  the  reacting  mole- 
cules. When  heat  is  evolved  the  change  is  exothermic, 

and  when  it  is  absorbed  the  change  is  endothermic.    All 

the  common  combustion  processes  are  exothermic  in  character  ; 
and  a  number  of  changes  will  be  met  with  in  metallurgical 
operations,  which  are  distinctly  endothermic  in  character. 
They  do  not  take  place  unless  heat  is  supplied  from  without. 
Energy  then,  as  we  know  it,  is  a  condition  of  matter,  and  like 
matter  can  undergo  a  variety  of  changes,  and,  in  any  self- 
contained  system  of  bodies,  the  amount  of  energy,  like  the 
amount  of  matter,  is  absolutely  fixed. 

MEASUREMENT  OF  HEAT. 

When  a  given  quantity  of  heat  enters  a  body  and  distributes 
itself  uniformly  through  the  substance  of  the  body,  a  change 
takes  place  in  its  thermal  properties,  and  the  temperature  of 
the  body  is  said  to  rise.  This  change  of  temperature  depends 


COMBUSTION  215 

upon  (a)  the  mass  of  matter  in  the  body,  and  (b)  its  power  of 
absorbing  heat.  The  latter  property  is  termed  its  specific 
heat,  and  is  dependent  upon  the  kind  of  matter  concerned  in 
the  absorption.  Specific  heat  and  temperature  are  regarded  as 
the  two  factors  of  heat  energy,  and  are  measured  by  reference 
to  arbitrary  standards.  The  temperature  of  a  body  indicates 
the  intensity  of  its  heat  energy,  and  is  the  factor  which 
determines  the  possibility  of  heat  entering  or  leaving  the  body 
under  given  conditions.  Thus  if  two  bodies  at  different 
temperatures  are  brought  together,  heat  will  pass  from  the 
body  at  the  higher  to  the  body  at  the  lower  temperature,  and 
the  exchange  will  go  on  until  the  intensity  of  the  heat  energy 
is  the  same  in  both.  The  temperature  of  the  one  body  rises 
and  that  of  the  other  falls,  until  it  is  the  same  for  both.  But 
the  quantity  of  heat  which  leaves  the  one  body  arid  enters  the 
other  also  depends  upon  their  specific  heats — i.e.,  upon  the 
other  factor  of  heat  energy.  Thus  if  a  number  of  bodies  at 
different  temperatures  are  brought  together  and  left  to  them- 
selves, exchange  of  heat  energy  will  take  place  between  them, 
and  they  will  finally  reach  one  dead  level  of  temperature,  the 
value  of  which  will  depend  upon  the  masses,  specific  heats,  and 
temperatures  of  all  the  bodies  in  the  system.  It  is  assumed  in 
this  general  statement  that  no  process  other  than  redistribution 
of  heat  energy  takes  place  in  the  system. 

Measurement  of  Temperature. — Advantage  is  taken  of 

the  fact  that  bodies  in  general  expand  when  heated  and  con- 
tract when  cooled,  and  that  this  expansion  and  contraction 
may  within  certain  limits  be  taken  as  a  measure  of  the  rise 
and  fall  of  their  temperature,  and,  therefore,  of  that  of  other 
bodies  with  which  they  are  in  contact.  The  instruments  used 
for  registering  moderate  temperatures  are  called  thermometers, 
and  those  for  high  temperatures  pyrometers.  The  mercurial 
thermometer  consists  of  a  glass  bulb  and  tube  nearly  filled 
with  mercury.  The  relation  between  the  capacity  of  the  bulb 
and  the  diameter  of  the  bore  is  such  that  the  metal  does  not 


216  METALLURGICAL  CHEMISTRY 

quite  leave  the  bore  at  the  lowest,  and  does  not  quite  fill  it 
at  the  highest  temperature  the  instrument  is  intended  to 
register.  The  fixed  points  of  the  scale  are  obtained  by  first 
immersing  the  instrument  in  melting  ice,  and  then  in  the 
steam  from  water  boiling  at  normal  pressure.  The  point  on 
the  stem  to  which  the  mercury  recedes  when  exposed  to  the 
temperature  of  melting  ice  is  marked  as  the  freezing-point  of 
water,  for  ice  melts  and  water  freezes  at  the  same  constant 
temperature  ;  the  point  on  the  stem  to  which  the  metal 
expands  when  exposed  to  the  steam  is  marked  as  the  boiling- 
point  of  water,  for  the  steam  from  water  boiling  at  a  constant 
pressure  has  a  constant  temperature.  The  space  between  these 
points  is  divided  into  degrees,  which  are  marked  on  the  scale. 

On  the  Centigrade  scale  there  are  100°  between  the  freezing 
and  boiling  points.  Thus  the  freezing-point  is  marked  0°  C., 
and  the  boiling-point  100°  C.  The  scale  may,  however,  be 
graduated  above  and  below  these  points ;  the  lower  limit  is 
the  freezing-point  (  -  39°  C.),  and  the  upper  limit  the  boiling- 
point  (350°  C.)  of  the  mercury  itself.  But  when  the  tempera- 
ture of  the  metal  approaches  either  of  these  limits  the  indica- 
tions are  uncertain,  and  it  may  be  said  generally  that  the  rate 
of  expansion  of  a  body  is  not  strictly  proportional  to  its 
increase  in  temperature,  and  the  hotter  it  gets  the  greater  is 
its  deviation,  so  that  in  accurate  instruments  corrections  have 
to  be  made  for  this  deviation.  In  the  case  of  mercury  the 
deviation  is  small  between  0°  and  100°,  and  can  be  corrected 
for  at  higher  temperatures,  so  that  some  mercurial  thermo- 
meters are  constructed  to  register  upwards  of  300°  C. 

The  temperature  scale  commonly  used  in  the  British  Isles 
is  the  Fahrenheit,  on  which  the  freezing-point  is  marked  32°, 
and  the  boiling-point  212°,  thus  giving  180°  between  the  two 
points.  The  conversion  of  a  reading  on  the  Fahrenheit  scale 
to  the  corresponding  reading  on  the  Centigrade  is  easily 

"Hi          -t  o  f\         f\  f\  pr 

effected,  for  ^  =  VQQ  =  -= ;  therefore  F  =  «  C,  and  C  =  g  F  ;  but 
in  converting  F.  into  C.  the  32°  below  the  freezing-point  must 


COMBUSTION  217 

be  subtracted  before,  and  in  the  conversion  of  C.  into  F.,  added 
after,  the  calculation  is  effected.  This  is  given  by  the  expres- 
sions : 

C  =  (F-32)J;F  =  f  C  +  32. 

To  determine  the  temperature  of  a  given  body,  the  ther- 
mometer is  brought  into  contact  with  it,  the  two  allowed  to 
come  into  equilibrium,  and  the  reading  then  taken.  The 
mass  of  the  thermometer  is  kept  as  small  as  possible,  so  that 
the  quantity  of  heat  entering  or  leaving  it  while  coming  into 
equilibrium  with  the  body  may  not  sensibly  interfere  with 
the  temperature  to  be  measured. 

The  construction  of  pyrometers  for  the  measurement  of  high 
temperatures  is  based  upon  various  principles,  such  as  the 
expansion  of  solids  and  gases ;  the  increase  in  the  electrical 
resistance  of  a  metal  wire  when  heated ;  the  current  generated 
by  a  thermo-electric  junction  when  its  temperature  is  in- 
creased, etc. ;  but  standard  works  on  metallurgy,  heat,  and  elec- 
tricity must  be  consulted  for  a  description  of  such  instruments. 

Specific  Heat. — It  is  a  matter  of  common  experience  that 
different  weights  of  the  same  substance  at  a  given  tempera- 
ture absorb  or  give  out  different  quantities  of  heat  when  they 
are  raised  or  lowered  through  the  same  range  of  temperature, 
and  that  these  quantities  are  strictly  proportional  to  the 
masses  of  the  substance  under  consideration.  Thus,  2  pounds 
of  water  at  50°  C.  would  liberate  twice  as  much  heat  in  falling 
to  20°  C.  as  1  pound  would  under  the  same  conditions.  Now, 
if  the  heat  given  out  by  a  given  weight  of  water  in  cooling 
through  a  given  range  of  temperature  is  absorbed  by  a  given 
weight  of  iron,  it  is  found  that  3*57  times  that  weight  of  lead 
is  required  to  absorb  the  same  quantity  of  heat ;  or  a  quantity 
of  heat  which  would  raise  100  grams  of  iron  through,  say,  10°  C. 
would  raise  357  grams  of  lead  through  the  same  temperature. 
This  difference  in  the  absorbing  power  of  the  two  metals  is 
said  to  be  due  to  the  difference  in  their  specific  heats.  The 


218  METALLURGICAL  CHEMISTKY 

heat-absorbing  power  of  water  is  greater  than  that  of  any 
other  compound  or  element,  and  is  taken  as  the  standard. 
The  specific  heat  of  water  =  1,  so  that  the  specific  heats  of  all 
other  pure  substances  is  less  than  unity  when  compared  with 
this  standard. 

The  method  of  determining  the  specific  heats  of  solid  bodies 
such  as  metals  may  be  thus  described :  A  known  weight  of 
water,  W,  is  placed  in  a  vessel  called  a  calorimeter,  and  its 
temperature,  T°,  taken  ;  a  known  weight  of  the  body  w,  which 
has  been  raised  to  a  temperature,  f,  is  then  immersed  in  the 
water  and  moved  about  until  the  two  come  to  the  same 
temperature,  and  the  temperature  of  the  water,  Tj0,  a'gain 
taken.  Now,  while  the  temperature  of  the  water  rises  from 
T°  to  Tj0,  the  temperature  of  the  body  falls  from  t°  to  T^, 
and  the  heat  given  out  by  the  body  is  equal  to  that  absorbed 
by  the  water  ;  but  the  quantity  of  heat  given  out  by  the  body 
is  proportional  to  its  mass,  specific  heat,  and  fall  in  tempera- 
ture, and  the  quantity  absorbed  by  the  water  to  the  same 
factors  ;  but  these  quantities  of  heat  are  equal,  therefore 

(t°  -  T^)  x  w  x  s  =  (Ty-T0)  x  W  x  S, 

where  s  and  S  are  the  specific  heats  of  the  body  and  water. 
But  as  S  =  1 — i.e.,  the  specific  heat  of  water — s,  the  specific 
heat  of  the  body,  is  the  only  unknown  quantity,  and  its  value 
is  easily  found  from  the  equation. 

EXAMPLE. — A  lump  of  platinum  weighing  2  ounces  is  raised  to  a 
temperature  of  410°  C.,  and  is  then  plunged  into  20  ounces  of  water 
at  16°  C. ;  the  resulting  temperature  of  the  water  is  found  to  be 
17'3°  C.  What  is  the  specific  heat  of  platinum  ? 

«°—T1°=410- 17-3  =  892-7,  and  T1°-T°  =  17'3-16  =  1  3. 
392-7x2xs  =  l'3x20xl; 
1-3x20     nAOO 
°r     S  =  S9^x2=°-088- 

A  correction  has  to  be  made  for  the  heat  absorbed  by  the 
calorimeter,  but  this  and  other  points  in  connection  with  the 
accurate  determination  of  specific  heats  cannot  be  discussed 
here.  The  specific  heat  of  platinum  between  0°  C.  and  100°  C. 


COMBUSTION  219 

is  0*0323,  and  increases  slightly  at  higher  temperatures. 
The  definite  relation  between  the  specific  heats  and  atomic 
weights  of  metals  is  discussed  on  p.  128. 

The  unit  of  heat  quantity  is  the  calorie,  which  is  the  amount 
of  heat  required  to  raise  the  temperature  of  1  gram  of  water 
from  0°  C.  to  1°  C.  The  great  calorie  or  kilogram  unit,  which 
is  equal  to  1,000  calories,  is  often  used  for  practical  work; 
also  the  quantity  of  heat  required  to  raise  1  pound  of  water  from 
32°  F.  to  33°  F.,  or  from  0°  C.  to  1°  C.,  may  be  taken  as  the 
unit. 

The  definition  of  unit  heat  quantity  makes  it  evident  that 
the  measurement  of  heat  energy  depends  upon  the  specific 
heat  of  the  absorbing  body,  and  the  temperature  through 
which  it  is  raised,  so  that  specific  heat  and  temperature  are 
the  factors  to  be  considered  in  heat  measurement. 

When  a  solid  changes  to  the  liquid  state  or  a  liquid  to  the 
gaseous  state,  heat  is  absorbed  without  change  in  temperature, 
but  the  heat  reappears  when  the  inverse  change  takes  place. 
The  heat  absorbed  by  melting  ice  is  80  calories,  and  that 
absorbed  by  boiling  water  is  537  calories.  These  numbers  are 
termed  the  latent  heats  of  fusion  of  ice,  and  of  vaporization 
of  water,  for  they  represent  the  quantities  of  heat  used  up  in 
changing  1  gram  of  ice  at  0°  C.  to  water,  and  1  gram  of  water 
at  100°  C.  to  steam  at  the  same  temperatures. 

SUMMARY. 

Carbon  and  hydrogen  are  the  principal  combustible  bodies, 
and  air  is  the  important  supporter  of  combustion.  Carbon 
dioxide  and  water  are  formed.  The  rate  at  which  heat  is 
developed  depends  upon  the  rate  at  which  the  chemical 
changes  take  place.  A  rapid  current  of  heated  air,  or  pure 
oxygen,  is  used  when  a  high  local  temperature  is  required. 
Flame  is  the  general  result  of  burning  a  combustible  gas. 
The  heat  energy  of  combustion  is  the  measurable  equivalent  of 
the  energy  of  chemical  attraction  which  existed  between  the 


220  METALLUKGICAL  CHEMISTRY 

combustible  body  and  the  oxygen  of  the  air  used  in  its  com- 
bustion. This  is  generally  true  for  all  chemical  changes  by 
which  heat  is  developed.  An  exothermic  change  is  when  heat 
is  developed  as  a  result  of  the  change ;  and  an  endothermic 
change  is  when  heat  is  absorbed  as  the  change  takes  place.  In 
the  one  case  energy  is  given  out,  and  in  the  other  it  is 
absorbed  by  the  reacting  bodies.  Specific  heat  and  tempera- 
ture are  the  two  factors  of  heat  energy. 

QUESTIONS. 

1.  What  is  understood   by   the  ignition-point   of   a   com- 
bustible body  ? 

2.  Give  a  short  description  of  two  processes  for  the  prepara- 
tion of  oxygen. 

3.  What  has  been  omitted  from  the  equation :  C  +  02  =  C02  ? 

4.  Describe   and   explain   three  characteristic  experiments 
with  oxygen  gas. 

5.  How  would  you  prove  that  a  combustible  body  is  not 
destroyed  as  it  burns  away  1 

6.  Explain  the  term  "inversion  of  combustion." 

7.  Explain  how  you  would  localize  or  distribute  the  heat  to 
be  obtained  by  burning  a  given  sample  of  fuel. 

8.  Explain  the  terms  "flame,"  "luminosity,"  and   "flame 
structure." 

9.  Give  a  description  of  the  Bunsen  burner. 

10.  Give  a  short  account  of  the  principle  of  the  conservation 
of  energy. 

11.  Explain  the  terms  "temperature"  and  "specific  heat." 

12.  How  is  the  specific  heat  of  a  metal  determined  ? 


CHAPTER  XIII 
PHOSPHOEUS  AND  ITS  COMPOUNDS 

WHEN  bones  are  heated  strongly  out  of  contact  with  air,  in  a 
similar  manner  to  coal  or  wood  (p.  178),  liquid  and  gaseous 
matters  are  given  off,  and  a  black  residue  is  left  in  the  retort. 
This  is  bone-black,  and  is  used  for  various  purposes.  If  it 
is  completely  burnt  in  a  current  of  air,  a  white  residue  known 
as  bone-ash  is  obtained,  which  consists  very  largely  of  calcium 
phosphate,  Ca3(P04)2.  This  substance  is  used  in  the  metal- 
lurgical operation  of  cupellation  to  absorb  molten  oxide  of 
lead,  and  also  for  the  extraction  of  phosphorus.  Calcium 
phosphate  is  a  constituent  of  the  older  rocks,  and,  when  they 
are  broken  down  in  the  slow  process  of  soil  formation,  remains 
in  the  soil.  It  is  gradually  converted  into  a  soluble  form,  and 
is  then  extracted  by  growing  plants,  which  assimilate  it,  and 
pass  it  on  to  animals  using  them  for  food. 

Extraction  of  Phosphorus. — Bone-ash  when  digested 
with  sulphuric  acid  is  converted  into  calcium  sulphate  and 
phosphoric  acid.  This  acid  is  soluble  in  the  remaining  acid 
liquor,  but  the  calcium  sulphate  is  practically  insoluble,  and 
separates  out  in  the  solid  state.  After  the  acid  liquid  has  been 
run  off  it  is  evaporated  to  a  syrup  in  an  iron  vessel,  mixed 
with  charcoal  powder,  and  further  heated  to  a  moderately 
high  temperature.  The  dry  mixture  is  then  transferred  to 
clay  retorts  arranged  in  a  furnace,  and  there  very  strongly 
heated.  Phosphorus  vapour  distils  over,  and  is  collected  in 
vessels  containing  water.  The  crude  phosphorus  thus  obtained 

221 


222  METALLURGICAL  CHEMISTEY 

is  purified  and  cast  into  sticks  ready  for  use.    The  changes  are 
expressed  by  the  equations  : 

Ca3(P04)2  +  3H2S04  =  3CaS04  +  2H3PO4. 


Natural  phosphates  are  found  in  some  parts  of  the  world  in 
considerable  quantities.  Apatite  and  phosphorite,  both  of 
which  contain  calcium  phosphate,  are  the  most  plentiful,  and 
may  be  used  for  the  extraction  of  phosphorus. 

Electric  Smelting*.  —  Recent  improvements  in  the  electric 
furnace  have  made  it  possible  to  smelt  either  a  natural 
phosphate  or  bone-ash,  when  mixed  with  charcoal,  for  the 
direct  production  of  phosphorus.  The  preliminary  treatment 
with  sulphuric  acid  is  thus  avoided,  and  a  saving  effected. 
The  extremely  high  temperature  of  the  furnace  makes  the 
reduction  possible. 

Properties  Of  PhOSphOPUS.  —  Ordinary  phosphorus  is  an 
amber-yellow,  waxy-looking  solid,  which  melts  at  44°  C.,  and 
boils  at  290°  C.  It  takes  fire  very  readily,  and  should  as  a 
rule  be  handled  and  kept  under  water.  It  is  readily  dissolved 
by  carbon  bisulphide,  CS2.  When  this  solution  is  evaporated 
the  phosphorus  is  left  in  such  a  finely  divided  condition  that 
it  bursts  into  flame  spontaneously.  It  is  also  dissolved  by  a 
solution  of  caustic  soda  or  potash  ;  but  in  this  case  a  reaction 
takes  place  by  which  the  gas  phosphoretted  hydrogen,  PH3, 
is  formed  and  liberated.  This  gas  when  impure  bursts  into 
flame  directly  it  comes  into  contact  with  the  air.  It  is 
poisonous,  and  has  the  odour  of  decaying  fish. 

When  yellow  phosphorus  is  heated  for  some  time  to  a 
temperature  about  240°  to  250°  C.  out  of  contact  with  oxygen 
it  is  gradually  converted  into  red  phosphorus,  which  is  a 
dark-red  powder,  insoluble  in  carbon  bisulphide,  CS2,  and 
perfectly  safe  to  handle  ;  but  when  heated  in  contact  with  air 
it  bursts  into  flame.  It  is  therefore  a  very  different  body  in 
appearance  and  properties  from  the  yellow  variety,  although 


PHOSPHORUS  AND  ITS  COMPOUNDS  223 

it   is  the  same  chemical  substance.      The   two  are  allotropic 
modifications  of  the  element. 

Phosphorus  unites  readily  with  metals  to  form  phosphides, 
which  are  analogous  to  the  sulphides  in  general  composition  ; 
but  they  are  not  so  readily  obtained  in  the  pure  state.  Some 
oxides  when  heated  are  attacked  by  phosphorus,  and  the 
phosphide  of  the  metal  is  formed.  This  is  the  case  with 
red-hot  lime.  Two  important  commercial  materials  are  phos- 
phor tin  and  phosphor  copper,  which  contain  phosphides  of 
the  metals.  When  lead  is  melted  with  an  excess  of  phosphorus, 
the  metal  absorbs  more  of  the  non-metal  than  it  can  hold 
when  in  the  solid  state.  This  excess  separates  as  crystals 
when  the  mass  solidifies.  The  lead  is  easily  dissolved  by 
dilute  nitric  acid,  and  the  phorphorus  left  behind  as  a  black 
residue.  This  resembles  red  phosphorus  closely,  except  in 
colour.  It  is,  in  fact,  another  allotropic  modification  of  the 
element.  It  may  be  noticed  here  that  the  association  of 
phosphorus  with  lead  brings  about  a  modification  of  the 
former  element ;  and  an  analogy  may  be  drawn  between  this 
change  and  the  conversion  of  amorphous  carbon  or  diamond 
carbon  to  the  graphitic  form  by  association  with  iron  under 
certain  conditions  (p.  167). 

Phosphorus  is  one  of  the  common  impurities  in  commercial 
iron,  and  passes  into  the  metal  from  phosphates  present  in  the 
ores,  and  reduced  by  carbon  at  the  high  temperature  of  the 
smelting  furnace  in  which  the  metal  is  extracted. 

When  phosphorus  burns  in  oxygen,  or  in  a  good  supply  of 
air,  phosphoric  oxide,  F2^5'  ^s  formed ;  but  when  the  air  is 
limited  and  the  combustion  slow,  a  large  proportion  of  the 
product  consists  of  lower  oxides,  of  which  phosphorous  oxide, 
P203,  is  the  most  important.  These  oxides  form  solutions 
which  have  a  sour  taste,  and  turn  litmus  red.  They  are, 
therefore,  acid-forming  oxides. 

Phosphoric  oxide  is  a  white,  powdery,  volatile  solid,  which 
rapidly  absorbs  moisture,  and  deliquesces  on  exposure  to  the 
air.  When  dissolved  in  cold  water  an  acid  is  formed,  which 


224  METALLURGICAL  CHEMISTRY 

unites  with  only  one  proportion  of  a  base,  and  forms  one  series 
of  salts.  It  resembles  nitric  acid  in  general  composition,  and 
is  known  as  metaphosphoric  acid,  HP03.  But  when  the 
solution  is  boiled  another  acid  is  formed,  which  is  capable  of 
giving  three  well-defined  series  of  salts.  It  is  tribasic  (p.  153), 
and  is  known  as  orthophosphoric  acid,  H3P04.  It  is  the 
normal  or  common  phosphoric  acid,  and  is  readily  prepared  by 
adding  red  phosphorus,  a  little  at  a  time,  to  warm  concentrated 
nitric  acid  as  long  as  the  solid  is  taken  up  The  solution  is 
then  evaporated  to  a  syrup,  which  is  the  practically  pure  acid. 
It  may  be  dissolved  in  water,  and  used  for  the  preparation 
of  phosphates.  The  common  prepared  salt  is  disodium 
hydrogen  phosphate,  Na2HP04,  12H20.  When  this  salt  is 
heated  it  loses  its  water  of  crystallization  first,  and  then  at  a 
higher  temperature  parts  with  another  molecule  of  water,  and 
is  converted  into  a  salt  of  pyrophosphoric  acid,  H4P207.  The 
last  change  is  expressed  thus  : 

2Na2HP04  =  Na4P207  +  H2O. 

Orthophosphate.    Pyropho;-phate. 

It  would  appear,  then,  that  phosphoric  oxide  combines  with 
water  in  three  proportions  to  form  three  distinct  acids,  thus  : 

1.  H20  +  P205  =  2HP03. 

Metaphos- 
phoric acid. 

2.  2H20  +  P205  =  H4P207. 

Pyrophos- 
phoric acid. 

3.  3H0  +  P0  =  2HP04. 


Orthophos- 
phoric acid. 

They  all  give  corresponding  salts  with  metallic  bases.  The 
natural  and  prepared  phosphates  are  for  the  most  part  ortho- 
salts.  The  salts  of  the  other  acids  are  not  often  met  with  in 
ordinary  work.  The  phosphates  of  the  alkaline  metals  and 
ammonium  are  soluble  in  water,  and  three  distinct  salts  of 
each  base  can  be  obtained,  if  necessary.  The  three  salts  of 
sodium  are  NaH2P04,  Na2HP04,  Na3P04.  Also,  it  is  possible 


PHOSPHORUS  AND  ITS  COMPOUNDS  225 

to  have  two  or  three  different  bases  in  the  same  salt.  Thus 
the  well-known  microcosmic  salt  has  the  formula  NaNH4HPO,t. 
It  is  sodium  ammonium  hydrogen  phosphate,  and  forms  well- 
defined  crystals. 

Phosphates  of  other  metals  are  for  the  most  part  insoluble 
in  water,  but  soluble  in  acids,  and  can  be  obtained  by  mixing 
together  solutions  of  a  soluble  phosphate  and  of  a  soluble  salt 
of  the  other  metal.  The  insoluble  phosphate  is  precipitated, 
and  can  be  separated  from  the  solution  by  filtration. 

Phosphorous  acid,  H3P03,  gives  a  series  of  salts  called  phos- 
phites, which  are  not  of  particular  importance. 

Phosphorus  unites  directly  with  chlorine,  bromine,  and 
iodine,  forming  important  compounds,  which  are  treated  fully 
in  advanced  works  on  chemistry. 

SUMMARY. 

Phosphorus  is  not  found  in  the  free  state  in  Nature.  Its 
principal  natural  compound  is  calcium  phosphate,  Ca3(P04)2. 
The  element  exists  in  at  least  two  distinct  allotropic  forms. 
It  combines  readily  with  metals  to  form  phosphides.  Phosphor 
tin  and  phosphor  copper  are  largely  used  in  the  preparation 
of  alloys.  The  principal  oxygen  compounds  of  the  element 
are  phosphoric  oxide,  P2O5,  and  phosphoric  acid,  H3P04. 
Phosphates  of  the  common  metals  are  easily  formed. 

QUESTIONS. 

1.  Give  a  short  description  of  the  usual  method  for  obtaining 
phosphorus. 

2.  Name  the  common  forms  of  phosphorus,  and  state  how 
they  differ  from  each  other. 

3.  What  happens  when  phosphorus  burns'  in  air  ? 

4.  Name  the  three  acids  corresponding  to  phosphoric  oxide. 
To  which  do  the  common  phosphates  belong1? 

5.  If  you  were  required  to  obtain  an  insoluble  phosphate  of 
a  metal,  how  would  you  proceed  1 

15 


CHAPTER  XIV 

SILICON  AND  ITS  COMPOUNDS  WITH  OXYGEN 
AND  METALS 

THE  element  silicon,  although  not  often  met  with  in  the  pure 
state,  is  present  in  abundance  in  the  earth's  crust,  combined 
with  oxygen  and  metals.  Its  most  important  compound  is 
the  dioxide,  silica,  Si02,  which  in  the  prepared  state  is  a 
perfectly  white  powder.  One  natural  form  of  this  oxide, 
quartz,  is  often  found  in  distinct  crystals,  which  are  sometimes 
as  transparent  as  glass.  It  is  an  acid-forming  oxide,  but  as  it 
is  quite  insoluble  in  water  under  ordinary  conditions,  this 
cannot  be  proved  in  the  usual  way.  It  is  also  very  refractory 
in  character,  requiring  a  very  high  temperature  to  soften  and 
fuse  it. 

Silica  is  not  decomposed  when  strongly  heated  with  carbon 
unless  a  metal  is  present  into  which  the  reduced  silicon  can 
pass  as  soon  as  it  is  liberated  from  the  oxide.  Thus,  when 
a  mixture  of  silica,  charcoal,  and  iron  is  very  strongly  heated, 
silicon  is  reduced  and  passes  into  the  metal.  This  is  the  source 
of  silicon  in  commercial  iron,  which  always  contains  that 
element,  though  sometimes  in  very  small  quantity.  On  the 
other  hand,  a  specially  prepared  iron  may  contain  as  much  as 
18  per  cent,  of  silicon.  The  non-metal  is  no  doubt  combined 
with  part  of  the  metal  in  the  form  of  iron  silicide,  which 
considerably  modifies  the  properties  of  the  remainder  of  the 
metal  with  which  it  is  associated.  When  iron  containing 
silicon  is  dissolved  in  an  acid,  and  the  solution  evaporated  to 
dryness,  a  residue  containing  insoluble  silica  is  obtained. 

226 


SILICON  AND  ITS  COMPOUNDS  227 

Silica,  when  finely  divided,  is  readily  dissolved  by  a  solution 
of  caustic  soda  or  potash,  and  an  alkaline  salt  of  silicic  acid 
is  formed.  With  caustic  soda  sodium  silicate,  Na4Si04,  is 
obtained  : 

4NaHO  +  Si02  =  Na4Si04  +  2H2O. 

A  similar  change  takes  place  when  any  form  of  silica  in  a 
finely-divided  state  is  fused  with  sodium  carbonate  in  a 
platinum  crucible.  The  reaction  proceeds  slowly,  and  carbon 
dioxide  is  liberated  : 

2Na2C03  +  Si02  =  Na4Si04  +  2C02. 

Fusion  mixture,  which  is  a  mixture  of  sodium  and  potassium 
carbonates,  is  more  effective  than  the  single  carbonate.  The 
fused  mass  can  be  extracted  with  hot  water,  which  dissolves 
out  the  soluble  silicate  together  with  the  excess  of  carbonate. 
When  dilute  hydrochloric  acid  is  added,  in  small  quantities  at 
a  time,  to  this  solution  the  excess  of  carbonate  is  decomposed 
without  the  silicic  acid  separating,  and  an  excess  of  hydrochloric 
may  be  thus  introduced.  This  decomposes  the  silicate  by 
converting  its  metal  into  a  chloride,  and  silica  in  combination 
with  water  is  retained  in  solution. 

If  this  acid  solution  is  poured  into  a  tray  fitted  with  a 
bottom  made  of  parchment  paper,  and  floating  on  water  in 
a  good-sized  vessel,  it  is  found  that  the  excess  of  hydrochloric 
acid,  together  with  the  whole  of  the  sodium  in  the  form  of 
chloride,  passes  through  the  parchment  into  the  water  under- 
neath, so  that  nothing  but  silica  is  left  in  solution  in  the 
tray.  There  is  reason  to  believe  that  the  silica  is  combined 
with  water  in  the  form  of  silicic  acid,  which  will  not  pass 
through  the  parchment.  But,  although  the  solution  can  be 
concentrated  considerably,  the  pure  acid  has  not  been  prepared. 
In  this  respect  it  may  be  compared  with  the  solution  of 
carbon  dioxide,  which  is  supposed  to  contain  carbonic  acid, 
H2C03. 

Graham,  who  discovered  and  investigated  the  properties 
of  bodies  taken  advantage  of  in  the  above  separation,  gave 

15—2 


228  METALLURGICAL  CHEMISTRY 

the  name  of  colloids  to  those  bodies  which,  when  in  solution, 
will  not  pass  through  the  parchment  paper,  and  the  name  of 
Crystalloids  to  those  which  will  so  pass.  He  also  gave  the 
name  of  dialysis  to  the  operation  of  separating  crystalloids 
and  colloids  when  together  in  solution.  Substances  resembling 
glue  are  colloids. 

Silicic  Acids. — There  is  evidence  in  favour  of  the  possible 
existence  of  two  silicic  acids :  orthosilicic  acid,  H4Si04,  and 
metasilicic  acid,  H2Si03.  It  is  readily  seen  that  the  first  is 
tetrabasic,  and  is  formed  by  the  combination  of  silica  with  2 
molecules  of  water ;  the  second  is  dibasic,  and  contains  only 
1  molecule  of  water. 

Silicates. — The  salts  of  sicilic  acid  are  very  numerous,  and 
are  found  in  great  variety  in  the  earth's  crust.  They  may 
be  formed  by  the  direct  combination  of  silica  with  basic 
oxides.  The  facility  with  which  a  particular  silicate  is 
formed  depends  upon  its  fusibility.  But  if  a  mixture  of 
silica  and  a  basic  oxide  in  the  proper  proportions  to  form  a 
certain  silicate  is  raised  to  or  a  little  above  the  melting-point 
of  the  silicate,  the  combination  will  take  place.  In  fact,  in 
some  cases  the  combination  takes  place  when  the  mixture  is 
exposed  to  a  temperature  below  its  fusing-point.  This  is  known 
as  fritting"  combination.  In  this  case  the  two  oxides  must 
be  very  finely  divided  and  intimately  mixed  together. 

EXP.  118. — Weigh  15  grams  of  lead  oxide,  and  mix  with  it  the 
necessary  weight  of  fine,  hard  sand  (silica  sand),  as  calculated  from 
the  equation  given  below.  Put  the  mixture  into  a  small  clay 
crucible,  and  raise  it  to  a  bright-red  heat  in  a  gas  muffle,  or  wind 
furnace.  When  thoroughly  melted,  pour  the  contents  of  the  crucible 
on  to  an  iron  plate.  The  resulting  silicate  when  cold  is  an  amber  - 
coloured  glass. 

The  two  oxides  unite  directly  at  a  moderate  temperature  to 
form  lead  silicate : 

2PbO         +         Si02     =2PbO.Si02. 
2(207  +  16) 


SILICON  AND  ITS  COMPOUNDS  229 

The  mode  of  writing  the  formula  of  a  silicate  adopted  in  the 
equation  is  the  one  most  used,  and  is  justified  by  the  way  in 
which  the  combination  takes  place. 

EXP.  119. — Mix  5  grains  of  finely-ground  red  copper  scale,  Cu20, 
with  the  necessary  quantity  of  silica  sand  to  form  the  silicate 
Cu.jO.Si02.  Heat  the  mixture  on  a  clay  dish  in  a  gas  muffle  to  a 
moderate  red  heat  for  half  an  hour.  Allow  the  dish  to  cool,  and 
grind  up  a  little  of  the  residue  in  a  mortar.  When  examined  with 
a  lens  the  particles  appear  to  be  all  alike.  The  particles  of  sand 
and  cuprous  oxide  fit  together,  and  the  combination  takes  place 
without  fusion. 

FUSIBILITY  OF  SILICATES. 

Single  Silicates. — Among  the  common  silicates  those  of 
the  alkaline  metals  and  lead  have  the  lowest  fusing-points  ; 
silicates  of  copper  and  iron  take  an  intermediate  place ; 
while  silicates  of  lime,  magnesia,  and  alumina  fuse  only  at 
high  temperatures.  But,  as  no  silicate  will  resist  the  excessive 
temperature  of  an  electric  furnace,  they  may  be  described 
generally  as  fusible  compounds  with  a  wide  range  of  fusing- 
points. 

A  normal  silicate  when  in  the  fused  state  will,  as  a  rule, 
dissolve  more  of  its  basic  oxide,  or  of  silica,  whichever  is 
presented  to  it.  In  the  first  case  the  normal  silicate  becomes 
a  basic  salt,  and  in  the  second  an  anhydro-acid  salt  (p.  154). 
The  addition  may  either  raise  or  lower  the  fusing-point 
of  the  mass.  There  is  no  general  rule  with  regard  to  this 
change. 

Complex  Silicates. — It  may  be  stated  generally  that  if 
two  chemical  compounds,  having  well-defined  melting-points, 
are  mixed  together  in  certain  proportions,  the  resulting 
mixture  will  have  a  lower  melting-point  than  either  of  the 
constituents  taken  singly.  Silicates  form  the  best  illustration 
of  this  general  principle,  and  every  advantage  is  taken  of  it  in 
metallurgical  operations.  For  example,  the  normal  silicates 
of  lime  and  alumina  are  separately  infusible  at  ordinary 
furnace  temperatures  ;  but  if  they  are  melted  together,  or 


230  METALLURGICAL  CHEMISTRY 

if  arrangements  are  made  for  their  simultaneous  formation  in 
certain  proportions,  a  readily  fusible  double  silicate  is  the 
result. 

The  formula  for  orthosilicic  acid  is  H4Si04,  or  2H2O.Si01 ; 
it  is,  therefore,  tetrabasic.  So  that  the  normal  salt  of  lime 
is  2CaO.Si02,  and  that  of  alumina  2Al2O3.3Si02.  Now,  either 
of  these  compounds  is  practically  infusible  at  ordinary  furnace 
temperatures,  but  if  they  are  heated  together  in  certain  pro- 
portions a  moderately  fusible  double  silicate  is  obtained. 
This  may  be  expressed  by  the  following  formula,  which  is 
usually  aimed  at  when  the  double  silicate  is  to  be  formed  in  a 
metallurgical  operation : 

6(2CaO.Si09)  +    2Al203.3Si02 
6(2x56  +  60)     2x102  +  3x60 

loST"  ~~384~~ 

The  proportions  given  are  easily  verified,  and  the  percentage 
of  each  constituent  can  be  calculated.  If  this  is  done,  mixtures 
of  the  oxides  are  readily  made  in  convenient  quantities  to 
give  approximately  the  above  composition. 

EXP.  120. — 1.  Mix  together  7  grams  of  quick-lime,  and  3'8  grams 
of  fine  sand  ;  put  the  mixture  into  a  carbon  crucible,  and  heat  it 
strongly  in  a  wind  furnace  for  twenty  minutes.  Put  a  lid  on  the 
pot,  and  place  it  well  down  in  the  fire,  so  as  to  expose  it  to  the 
highest  temperature  the  furnace  will  give.  Remove  the  pot  from 
the  fire  at  the  end  of  stipulated  time,  and  allow  it  to  cool.  2.  Mix 
together  2'1  grams  of  alumina  and  1/9  grams  of  fine  sand,  and 
follow  the  instructions  given  for  No.  1.  On  examination  the  mixtures 
are  found  to  be  fitted  together,  and  the  one  containing  the  lime 
may  show  signs  of  incipient  fusion.  3.  Grind  the  two  mixtures 
thoroughly  together  in  an  iron  mortar  ;  transfer  the  new  mixture 
so  obtained  to  the  carbon  crucible,  and  heat  it  again  for  twenty 
minutes  under  the  same  conditions  as  before.  Remove  the  crucible 
from  the  furnace,  allow  it  to  cool,  and  examine  the  result  of  the 
experiment. 

If  the  experiment  is  carefully  conducted,  a  button  of  the 
perfectly  fused  double  silicate  will  be  obtained.  It  is  very 
brittle  and  readily  fractured.  Very  often  it  breaks  up  while 
cooling,  which  is  due  to  unequal  contraction  in  different  parts 


SILICON  AND  ITS  COMPOUNDS  231 

of  the  mass.  The  appearance  of  the  fractured  surface  depends 
somewhat  upon  the  furnace  conditions  and  on  the  rate  of 
cooling.  It  may  vary  from  a  bright  glassy  to  a  dull  earthy 
appearance. 

A  similar  experiment  may  be  made  with  the  corresponding 
compounds  of  metasilicic  acid,  H2O.Si02.  The  composition 
of  the  most  fusible  double  silicate  of  lime  and  alumina  is  given 
by  the  formula  : 

3(CaO.Si02)  +  Al203.3Si02. 

The  metasilicate  of  lime  will  show  more  signs  of  fusion  than 
the  corresponding  orthosilicate ;  and,  if  the  temperature  is 
very  high,  the  metasilicate  of  alumina  will  soften  considerably. 
It  is,  therefore,  more  fusible  than  the  corresponding  ortho- 
silicate.  The  double  silicate  is  also  more  fusible,  and  forms 
the  normal  slag  of  furnaces  working  at  a  moderate  tempera- 
ture. Orthosilicate  slags  come  from  hotter  furnaces. 

If  a  small  quantity  of  a  basic  oxide  is  added  to  a  fused 
silicate,  it  is  dissolved,  and  exerts  an  influence  upon  (1)  the 
fusibility  of  the  silicate  and  (2)  its  appearance.  In  most  cases 
the  fusibility  is  increased,  but  this  is  not  always  so;  an 
addition  of  zinc  oxide,  ZnO,  for  example,  decreases  the 
fusibility.  Various  colours  are  imparted  to  silicates  by  the 
addition  of  oxides  of  iron  (green),  copper  (red),  cobalt  (blue), 
etc.  Oxide  of  tin  renders  the  silicate  opaque,  as  in  white 
enamels. 

Fireclay  is  an  acid  silicate  of  alumina  containing  com- 
bined water,  which,  together  with  the  fineness  of  its  particles, 
renders  the  clay  plastic  when  mixed  with  water.  Its  com- 
position may  be  represented  by  #Si02.2Al203.3Si02.2H20, 
when  #Si02  represents  a*variable  proportion  of  the  acid  oxide. 
This  material,  when  approximately  pure,  is  very  refractory ; 
but  the  addition  of  small  quantities  of  basic  oxides  increases 
its  fusibility  and  renders  it  unfit  for  fire-resisting  purposes. 
This  effect  depends  upon  the  nature  and  proportion  of  the 
added  oxide.  Soda,  potash,  and  oxide  of  lead  exert  the 


232  METALLUEGICAL  CHEMISTRY 

greatest  influence ;  lime  and  oxide  of  iron  are  also  very 
hurtful,  the  latter  especially  so  if  carbon  or  reducing  gases 
are  present  to  reduce  the  ferric  oxide,  Fe203,  to  ferrous  oxide, 
FeO.  These  oxides  unite  with  the  excess  of  silica,  and  fusible 
double  silicates  are  formed,  and  if  present  as  impurities  in  the 
clay,  are  injurious  to  its  refractory  character. 

Exp.  121. — Mix  a  few  grams  of  finely-powdered  fireclay  with 
5  per  cent,  of  dry  sodium  carbonate,  which  will  carry  about  3  per 
cent,  of  soda,  Na20,  into  the  mixture.  Heat  the  mixture  strongly 
in  a  carbon  crucible  for  twenty  minutes,  and  examine  the  result. 

Glass. — If  the  experiments  described  above  are  carried  out, 
the  tendency  of  fused  silicates  to  form  the  transparent  amor- 
phous body  known  as  glass  is  clearly  demonstrated.  Long 
experience  has  taught  the  glass-maker  the  mixtures  of  silicates 
which  give  the  best  results  for  particular  purposes.  Thus, 
Window  glass  is  a  double  silicate  of  lime  and  soda ;  Bohemian 
glass,  which  is  a  double  silicate  of  lime  and  potash,  will  with- 
stand a  moderately  high  temperature  without  softening,  and 
changes  in  temperature  without  cracking:  it  is,  therefore, 
very  useful  for  combustion-tube  and  for  chemical  apparatus 
generally.  Flint  glass  is  a  double  silicate  of  lime  and  lead 
oxide,  and  softens  readily  in  an  ordinary  gas  flame ;  it  is  used 
for  making  glass  tube  which  is  required  to  be  easily  bent. 
Glass  has  to  be  carefully  annealed  by  allowing  it  to  cool  very 
slowly  after  it  has  been  blown  into  shape.  The  plastic  state 
of  hot  glass,  which  extends  through  a  wide  range  of  tempera- 
ture, renders  it  possible  for  the  glass-blower  to  do  with  his 
material  what  the  potter  does  with  plastic  clay.  The  intro- 
duction of  small  quantities  of  metallic  oxides  brings  about  the 
variations  in  colour  observed  in  ornamental  glasses ;  and  a 
very  small  quantity  of  free  metal  in  an  extremely  finely  divided 
state  may  influence  the  colour,  as  is  the  case  writh  gold  in  ruby 
glass.  The  opalescence  in  some  varieties  of  glass  is  caused  by 
the  added  oxides  not  completely  dissolving  in  the  molten  mass, 
and  the  very  fine  solid  particles  being  suspended  in  the  fluid 
matter  while  it  sets. 


SILICON  AND  ITS  COMPOUNDS 


233 


Nomenclature  of  Silicates. — A  system  of  naming,  which  has 
been  very  largely  adopted  by  metallurgists,  depends  upon  the 
proportion  between  the  number  of  atoms  of  oxygen  in  the  acid 
and  basic  portions  of  the  given  silicate.  This  is  clearly  depen- 
dent upon  the  assumption  that  a  silicate  is  a  combination  of  two 
or  more  oxides,  one  of  them  always  being  silica.  Take,  for 
example,  the  sodium  salt  of  orthosilicic  acid,  2Na2O.Si02.  The 
relation  here  is  1  :  1,  for  there  are  2  atoms  of  oxygen  in  each 
part  of  the  compound.  But  with  the  sodium  metasilicate, 
Na2O.Si02,  the  relation  is  2  :  1,  for  there  are  2  atoms  of 
oxygen  in  the  acid  portion  to  1  atom  in  the  basic  portion. 
The  number  of  atoms  of  the  metal,  which  varies  with  its 
valency,  does  not  influence  the  character  of  the  silicate.  The 
ratio  1  :  1  gives  rise  to  monosilicates,  and  the  ratio  2  :  1 
bisilicates.  But  this  does  not  account  for  the  anhydro-acid 
silicates,  of  which  there  are  numerous  examples.  Thus  in  the 
trisilicates  the  ratio  is  3  :  1,  and  in  the  sesquisilicates  it  is 
3:2.  There  are  also  silicates  which  contain  an  excess  of  the 
basic  oxide,  and  the  most  important  of  these  are  included  in 
the  subsilicates  in  which  the  ratio  is  1  :  2. 


TABLE  OF  SILICATES. 


RATIO  OF   NUMBER  OF 

ATOMS  OFOxYGEM  IN 
ACID  OXIDE  TO  NUM- 
BER   OF    ATOMS     OF 

MONOVALENT 
METAL, 

DIVALENT  METAL, 
M"O. 

TRIVALENT  METAL, 
M"'.203. 

OXYGEN      IN      BASIC 

OXIDE. 

1  :  2 

Sab           4Na2O.Si02 

4CaO.Si02 

4ALOs.3Si09 

1  :  1 

Mono 

2Na2O.Si02 

2CaO.Si02 

2AUOs.3SiO., 

2  :  1 

Bi 

Na2O.Si02 

CaO.SiO., 

Al20,.3SiO., 

3  :  1 

Tri 

2Na2O.Si02 

2Ca0.3Si02 

2Alo63.9Si(X 

q    .    Q 

3:2 

Sesqui 

4Na.20.3Si02 

4Ca0.3Si02 

4Al,03.9Si02 

The  metals  commonly  present  in  silicates  are  :  (1)  Mono- 
valent,  general  formula  of  oxide  M'20,  where  M  represents  an 
atom  of  the  metal ;  (2)  divalent,  M"0 ;  and  (3)  trivalent, 


234  METALLURGICAL  CHEMISTEY 

M'"2O3.  Often  all  three  of  these  types  of  oxides  are  present 
in  the  same  well-defined  silicate.  It  is  the  readiness  with 
which  acid,  basic,  and  complex  silicates  seem  to  form  that 
makes  these  bodies  so  numerous. 

When  a  natural  or  prepared  silicate  is  in  the  form  of  a 
glassy  or  stony-looking  body,  it  is  impossible  to  state  for  certain 
that  it  is  a  well-defined  compound,  although  its  composition 
can  be  accurately  determined.  But  when  it  is  found  or 
obtained  in  well-defined  crystals,  this  difficulty  disappears,  and 
an  exact  formula  can  be  assigned  to  it.  Such  bodies  often 
occur  in  the  earth's  crust,  and  may  also  be  prepared  artificially. 
In  this  respect  it  is  to  be  borne  in  mind  that  slow  cooling  of  a 
molten  mass  favours  the  formation  of  crystalline  structure 
and  the  development  of  individual  crystals. 

Mono-,  di-,  and  tri-valent  metals  are  often  associated  in 
the  same  natural  silicate,  and  empirical  formulae,  as  simple  as 
possible,  are  commonly  used  in  such  cases.  Thus  the  well- 
known  mineral  felspar  is  written  empirically  KAlSi308,  and 
this  tells  the  beginner  very  little,  except  that  it  is  a  double 
silicate  of  potassium  and  aluminium ;  but  if  it  is  noticed  that 
the  monovalent  K'  and  trivalent  Al"'  give  a  tetravalent  com- 
bination (KA1)"",  which  is  equivalent  to  2  atoms  of  oxygen, 
the  formula  may  be  written  (KAl)02.3Si02,  and  the  compound 
is  seen  to  be  a  trisilicate.  Also,  potassium  may  be  more  or 
less  completely  replaced  by  sodium  without  altering  the 
crystalline  form.  In  this  case  the  different  compounds  are 
isomorphous  (p.  128).  Isomorphism  plays  a  very  important 
part  in  the  mineral  world.  Na20  and  K20 ;  CuO,  MgO,  MnO, 
and  FeO ;  A12O5,  Fe203,  and  MiiaOg  replace  each  other  without 
altering  the  crystalline  form  of  the  minerals  in  which  they 
occur. 

These  natural  bodies  of  definite  composition  are  the 
minerals  which  help  to  form  the  various  rocks  in  the  earth's 
crust.  Thus,  granite  is  made  up  of  three  well-defined  minerals  : 
felspar,  mica,  and  quartz,  and  the  crystals  of  these  bodies 
were  formed  during  the  extremely  slow  cooling  of  the  molten 


SILICON  AND  ITS  COMPOUNDS  235 

rock.  Granite  is  readily  melted  in  an  ordinary  wind  furnace, 
but  does  not  recover  its  crystalline  structure  on  cooling. 
The  solidified  mass  resembles  the  slags  formed  in  smelting 
operations. 

SUMMARY. 

Silica,  Si02,  the  only  oxide  of  silicon,  is  the  most  generally 
occurring  compound  in  the  earth's  crust.  Quartz,  flint,  and 
sand  are  its  commonest  forms.  In  combination  with  metallic 
oxides  it  forms  the  natural  and  prepared  silicates,  which  are 
a  very  numerous  class  of  bodies.  There  are  two  silicic  acids, 
which  give  rise  to  two  well-defined  series  of  compounds, 
orthosilicates  and  metasilicates.  These  normal  silicates, 
when  in  the  molten  state,  readily  take  up  an  excess  of  either 
silica  or  metallic  oxides,  and  form  basic  and  anhydro-acid 
silicates.  Different  silicates  also  mix  together  readily,  and  a 
large  number  of  complex  bodies  are  thus  formed.  As  a  rule, 
a  complex  silicate  is  more  fusible  than  a  single  silicate.  The 
element  silicon  unites  with  metals  to  form  silicides.  It  is 
commonly  found  as  an  impurity  in  commercial  iron. 

QUESTIONS. 

1.  Give  a  short  description  of  the  physical  and  chemical 
properties  of  silica. 

2.  What  is  the  general  effect  of  mixing  silicates  together, 
upon  the  f using-point  of  the  mixture  1 

3.  Describe  an  experiment  in  which  a  silicate  is  formed. 

4.  What  is  fireclay,  and  to  what  does  it  owe  its  plasticity  1 

5.  How  are  silicates  classified  for  metallurgical  purposes  ? 

6.  What  do  you  understand  by  the  isomorphous  substitution 
of  one  metallic  oxide  for  another  in  a  silicate  1 


CHAPTER  XV 
WEIGHTS,  MEASUEES,  AND  APPARATUS 

THE  Metric  System  of  weights  and  measures,  which  is  of 
French  origin,  is  almost  universally  used  for  scientific  purposes. 
It  is  a  decimal  system,  and  on  that  account  is  very  easy  to 
work  with. 

Length. — The  unit  of  length  is  the  metre,  which  is 
divided  into  10  parts,  or  decimetres;  into  100  parts,  or 
Centimetres;  and  into  1,000  parts,  or  millimetres.  The 
Latin  prefixes,  deci-,  centi-,  and  milli-,  mean  ^  T^y,  and 
unrip  respectively,  and  denote  the  submultiples  of  the 
metre. 

The  Greek  prefixes,  deka-,  hecto-,  and  kilo-,  denote  10,  100, 
and  1,000,  respectively,  and  are  used  for  multiples  of  the  metre. 
Thus,  1  kilometre  =  1,000  metres. 

Volume. — The  common  •  measure  of  volume  is  the  litre. 
Its  capacity  is  equal  to  that  of  a  cube,  1  decimetre,  or  10 
centimetres  on  the  side.  The  litre  =  1  cubic  decimetre  =  1,000 
cubic  centimetres.  Any  volume  less  than  a  litre  is  usually 
expressed  in  cubic  centimetres.  Thus,  J  litre  =  500  cubic 
centimetres  =  500  c.c.,  c.c.  being  the  contraction  for  cubic 
centimetres. 

Weight. — The  unit  of  weight  is  derived  from  the  metre  ; 
it  is  the  weight  of  a  cubic  centimetre  of  pure  water  at  its 
point  of  maximum  density  (4°  C.),  and  is  called  the  gramme 
or  gram.  The  submultiples  are  the  decigram,  or  y1^  gram ; 
the  centigram,  or  yj^  gram  ;  and  the  milligram,  or 

236 


WEIGHTS,  MEASUEES,  AND  APPAKATUS         237 

gram.  The  Greek  prefixes  are  used  for  the  multiples.  Thus, 
the  kilogram  =  1,000  grams. 

ENGLISH  EQUIVALENTS  OF  THE  METRIC  UNITS. 

The  metre      =  39'37  inches. 
,,     litre         =    1*76  pints. 
„     gram        =  15'432  grains. 
,,     kilogram  =    2*2046  pounds  avoirdupois. 

The  equivalents  of  the  multiples  or  submultiples  are  easily 
obtained  by  moving  the  decimal  point  to  the  right  for 
multiples,  and  to  the  left  for  submultiples. 

Thus,  1  kilometre  =39,370  inches. 
1  centigram  =  0*1 5342  grain. 

The  following  will  be  found  useful  in  making  comparative 
measurements  : 

1  pound  avoirdupois  =  7,000  grains. 

i  ounce  ,,          =437 *5       „ 

1  pound        troy        =5,760       „ 

1  ounce  ,,          =     480       ,, 

1  pint  =       20  fluid  ounces. 

1  pint  of  water  at  15'5°  C.  weighs  20  ounces  avoirdupois. 
1  gallon  of  water  at  15 '5°  C.  weighs  10  pounds  avoirdupois. 
1  cubic  foot  of  water  at  15'5°  C.  weighs  62-4  pounds  avoirdu- 
pois nearly. 

The  Balance. — The  most  casual  observer  is  familiar  with 
the  ordinary  beam  scales  used  in  everyday  life ;  and,  as  the 
more  sensitive  balance  used  for  scientific  purposes  is  exactly 
the  same  in  principle,  no  difficulty  should  be  experienced  in 
using  it.  Accuracy  in  weighing  is  only  a  matter  of  care  and 
experience. 

The  beginner  should  examine  the  balance  he  is  using,  and 
notice  that  it  consists  of  an  accurately  made  beam,  supported 
by  knife  edges  which  rest  on  smooth  horizonal  plates  or  in 


METALLURGICAL  CHEMISTRY 


V-shaped  grooves,  and  having  two  scale-pans  suspended  in  a 
similar  manner,  one  from  each  end.  The  points  of  suspension 
of  the  pans  are  at  exactly  the  same  distance  from  the  point 
of  suspension  of  the  beam,  and  the  two  halves  of  the  beam 

are  exactly  similar  to  each 
other.  When  the  balance 
is  properly  set  the  pointer 
should  be  perfectly  vertical 
if  the  beam  is  at  rest,  or 
should  swing  to  nearly 
equal  distances  on  each 
side  of  the  vertical  division 
on  the  scale  at  the  foot  of 
the  supporting  pillar  if  it 
is  in  motion.  The  smaller 
the  weight  required  to  be 
added  to  either  one  side  or  the  other  in  order  to  disturb  this 
equality,  the  more  sensitive  the  balance  is,  and  the  more 
accurate  the  weighing  which  can  be  made  upon  it.  A  cheap 
and  sufficiently  accurate  form  of  balance  is  shown  in  Fig.  41. 

The  Weights. — The  Metric  System  of  weights  is  almost 
universally  used  for  scientific  purposes,  and  a  box  of  gram 
weights  containing  from  100  grams  to  1  milligram  will  be 
found  sufficient  for  most  purposes.  The  individual  weights — 
— are  commonly  marked  as  shown  in  the  diagram. 


FIG.  41. 


Gramsj 

Decigrams 

Centigrams 

Milligrams 


100 
10 
0-5 
0-05 


50 
5 

0-2 
0-02 


20 
2 

0-1 
0-01 


20 
1 

0-1 
0-01 


0-005  0-002  0-001  0-001 


Sometimes  it  is  found  convenient  to  use  grain  weights,  and 
sets  varying  from  1,000  grains  to  0*5  grain  would  be  found 
sufficient  for  the  work  described  in  this  book.  In  converting 
the  weight  in  grams  given  in  the  text  into  grains  it  will  be 


WEIGHTS,  MEASURES,  AND  APPARATUS          239 

sufficient  to  multiply  the  weight  in  grams  by  15*5.     1  gram  = 
15-432  grains. 

Weighing. — The  body  to  be  weighed  is  usually  placed  in 
the  left-hand  pan,  the  weights  in  the  right-hand  one.  The 
weights  should  be  used  systematically,  as  weighings  are  made 
more  rapidly  by  so  doing.  The  beam  can  be  raised  or  lowered 
by  means  of  a  lever  arrangement  worked  from  the  front  of  the 
stand,  and  bodies  to  be  weighed  should  not  be  put  into  the 
pan,  nor  should  weights  be  added  while  the  beam  is  swinging 
freely.  When  weighing  it  is  not  necessary  to  wait  for  the  beam 
to  come  to  rest,  for  when  the  pointer  swings  to  very  nearly 
equal  distances  on  each  side  of  the  perpendicular  mark  on 
the  graduated  scale  in  front  of  which  it  moves,  the  weighing 
is  finished.  Bodies  which  are  at  all  likely  to  soil  the  scale-pan 
should  be  weighed  in  a  counterpoised  watch-glass  or  crucible. 
In  fact,  it  is  a  good  plan  to  make  all  weighings  in  this  way,  as 
any  little  inaccuracy  in  the  balance  is  thus  neutralized.  For 
example,  if  a  watch-glass  is  weighed,  and  a  body  then  weighed 
in  it,  the  difference  between  the  two  weights  is  the  weight  of 
the  body.  The  beginner  is  strongly  advised  to  get  into  the 
habit  of  writing  down  the  results  of  weighings  before  he  leaves 
the  balance.  A  little  practice,  especially  with  tho  assistance 
of  a  teacher,  will  enable  him  to  weigh  accurately,  and  soon 
convince  him  that  methodical  weighing  is  much  more  rapid 
than  an  erratic  putting  on  and  pulling  off  of  weights. 

Measuring"  Vessels. — The  volume  measures  of  the  Metric 
System*  are  almost  always  used  in  volumetric  work.  The 
cubic  decimetre  or  litre  is  the  most  important,  and  contains 
1,000  cubic  centimetres.  In  the  experiments  described  in 
the  text  the  volumes  used  are  less  than  a  litre,  and  are 
generally  expressed  in  cubic  centimetres. 

The  burette  is  a  long,  narrow,  glass  tube,  open  at  one  end 
and  fitted  at  the  other  with  a  tap  arrangement  through  which 

*  See  the  table  of  the  Metric  System  of  weights  and  measures  at  the 
beginning  of  the  chapter. 


240 


METALLUKGICAL  CHEMISTRY 


small  quantities  of  the  contained  liquid  can  be  run.  The  tap 
must  be  so  under  control  that  the  liquid  may  be  made  to 
leave  the  burette  one  drop  at  a  time,  if  necessary.  In  cheap 
burettes  it  is  usual  to  connect  the  body  with  the  nozzle  by 
means  of  a  piece  of  rubber  tube,  and  to  close  this  by  a  spring 
pinch-tap.  A  glass  tap  is  shown  in  Fig.  42.  A  very  simple 
and  effective  substitute  for  the  pinch-tap  may  be  made  by 
introducing  a  piece  of  glass  rod  J  inch  long,  rounded  at  the 
ends,  and  of  such  diameter  as  to  require  gently  forcing  into 
the  rubber  tube  connection  between  the  nozzle  and  the  body 
of  the  burette.  This  completely  closes  the 
burette  while  at  rest,  but  when  the  rubber 
tube  just  over  the  rod  is  squeezed  between  the 
finger  and  thumb  a  channel  in  the  rubber  is 
made  down  the  side  of  the  rod  by  which  the 
liquid  can  escape.  On  removing  the  pressure 
the  burette  is  at  once  closed.  The  body  of 
the  burette  is  graduated  between  A  and  B 
into  fifty  divisions,  each  equal  to  1  cubic  centi- 
metre, so  that  if  a  liquid  is  level  with  the  zero 
mark  at  A,  and  is  run  out  until  it  is  level  with 
B,  exactly  50  cubic  centimetres  are  delivered. 
If  the  level  is  reduced  to  some  point  between 
A  and  B,  then  the  graduation  mark  at  that 
point  gives  the  volume  delivered.  The  sur- 
face of  the  liquid  in  the  burette  is  curved,  and  the  read- 
ings are  best  made  by  reference  to  the  lower  or  convex 
surface  of  the  liquid.  The  burettes  usually  sold  are  not  very 
accurate ;  but  if  any  connected  set  of  measurements  are  made 
with  the  same  burette  a  slight  inaccuracy  does  not  appreciably 
affect  the  results.  For  liquids  which  corrode  rubber  tube, 
burettes  with  glass  taps  must  be  used. 

Measuring1  Flasks  are  usually  sold  in  four  sizes— i.e., 
1,000  c.c.,  500  c.c.,  250  c.c.,  and  100  c.c.  They  sometimes 
have  two  graduation  marks  on  the  neck,  one  to  contain  and 


FIG.  42. 


WEIGHTS,  MEASURES,  AND  APPARATUS         241 


FIG.  43. 


one  to  deliver.     As  these  flasks  expand  or  contract  as  they 

are  heated  or  cooled,  the  reading  is  only  strictly  true  at  the 

temperature    at    which 

the    graduation    was 

made.     This  is  usually 

the  normal  temperature 

of    a    laboratory,     i.e., 

15°C.    Thus,  if  a  liquid 

is  poured  into  the  dry 

flask    until     the    level 

reaches  the  lower  gradu- 
ation mark  />,  Fig.  43, 

the    flask   contains   the 

registered     volume     of 

liquid ;  but  if  the  level 

is  brought  to  the  upper  graduation  mark  a,  then  it  will  deliver 

the  registered  quantity  if  the  liquid  is  poured  into  another 

vessel.     The  extra  liquid  between  the  two  marks  compensates 

for  that  which  adheres  to  the  sides  of  the  flask  after  the  main 

bulk  has  been  poured  out. 

Three  very  useful  measures  are  shown  in  Fig.  44,  which  may 

be  used  for  rough  measurements.     They  deliver  500,  250,  and 

50  c.c.  respectively.  For  the 
rapid  measurement  of  a  few 
cubic  centimetres  of  liquid  a 
pipette  can  be  made  by  draw- 
ing off  to  a  point  in  the 
gas  flame  a  piece  of  glass 
tube  about  T8F  of  an  inch 
internal  diameter  and  8  inches 
long.  To  graduate  it  into 
cubic  centimetres,  dip  the 
narrow  end  into  water,  and 
when  some  of  the  liquid  has 

entered,  remove  the  tube  and  let  out  as  much  of  the  liquid 

as    will   run    from   the    narrrow    end.      Without   disturbing 

16 


c.c. 

p 

=-500 

1-400 

1-300 

c.c. 

p 

H-300 

"      201 

c.c. 

p 

=- 

rr90 

1-200 

^-150 

—  80 

—  70 



—  6C 

__ 

10( 

—50 

=-100 

EE-JO 

—40 
—  3C 
-20 

^ 

=i  

S  r 

^  ^ 

—  10 

s 

FIG.  44. 


242  METALLURGICAL  CHEMISTRY 

the  small  quantity  of  liquid  not  delivered,  close  the  narrow 
end  with  a  little  stiff  grease,  and  run  in  at  the  open  end  1  c.c. 
of  water  from  a  burette ;  mark  the  level  of  the  liquid  by 
scratching  the  side  of  the  tube  with  a  file ;  run  in  a  second  cubic 
centimetre  without  disturbing  the  first,  and  mark  as  before ; 
continue  this  until  about  two-thirds  of  the  tube  has  been 
graduated.  The  pipette  will  then  deliver  the  number  of  cubic 
centimetres  marked  upon  it.  To  use  it,  place  the  tube  in  the 
liquid  of  which  the  measured  quantity  is  required ;  the  liquid 
enters  the  tube,  and  when  the  required  quantity  is  in,  place  the 
finger  on  the  top  of  the  tube  and  remove  it.  To  deliver  the 
liquid,  remove  the  finger,  and  it  will  run  out.  A  very  useful 
pipette  made  in  various  sizes  is  shown  in  Fig.  44. 

Wash-bottle. — A  very  convenient  form  of  wash-bottle  for 
quantitative  work  is  shown  in  Fig.  45.  It  consists  of  a  glass 
flask,  A,  fitted  up  as  shown  in  the  figure. 
The  nozzle,  B,  is  attached  by  means  of  a 
piece  of  rubber  tube,  and  the  neck,  C,  is 
bound  with  a  layer  of  thick  string.  The 
flexible  nozzle  allows  of  a  stream  of  water 
being  sent  in  any  required  direction.  If 
the  flask  is  thin  the  water  in  it  may  be 
raised  to  boiling,  either  on  the  sand-bath  or 
on  wire  gauze,  when  the  necessity  of  wash- 
ing with  hot  water  arises. 

45^  Other   Apparatus.  —  In    devising    the 

experiments  described  in  the  text,  the  chief 
aim  has  been  to  use  such  apparatus  as  the  student  in  an 
elementary  laboratory  may  be  expected  either  to  possess 
or  to  have  access  to.  Expensive  apparatus  of  the  easily- 
broken  type  has  been  either  excluded  or  attempts  have 
been  made  to  imitate  it  in  a  simple  manner.  When  more 
accurate  apparatus  than  that  described  can  be  used,  better 
results  may  be  expected.  But  the  student  often  gains  some 
knowledge  by  fitting  up  a  piece  of  apparatus  consisting 


WEIGHTS,  MEASUEES,  AND  APPARATUS         243 

of  several  parts,  which  by  the  ingenuity  of  the  glass-blower 
may  be  obtained  complete.  The  principle  involved  in  the 
saying,  "  You  press  the  button,  we  do  the  rest,"  is  no  doubt 
useful  for  many  purposes,  but  it  is  bad  from  an  educational 
point  of  view. 

The  table  furnace  shown  in  Fig.  8  was  evolved  from  a 
simpler  form  consisting  of  an  iron  stand  carrying  some  loose 
fireclay  slabs,  and  having  two  ordinary  Bunsen  burners 
underneath.  In  the  first  modification  the  movable  burners 
were  replaced  by  three  fixed  burners,  and  in  the  last  the 
fireclay  slabs  were  replaced  by  a  rectangular  fireclay  chamber, 
surmounted  by  a  sheet-iron  tube  to  increase  the  draught. 
The  temperature  obtained  in  this  chamber  is  very  high  for 
a  table  furnace,  and  experiments  which  in  the  ordinary  form 
of  furnace  either  fail  or  are  only  partially  successful  can  be 
carried  out  with  it. 

Conical  flasks  are  much  more  convenient  for  general  work 
than  the  ordinary  globular  variety.  The  8,  12,  and  16  ounce 
sizes  are  the  most  useful. 

Porcelain  Crucibles.  —  The  best  Berlin  crucibles  are  the 
cheapest  in  the  end,  and  Nos.  1  and  2  are  the  useful  sizes. 

Combustion-tube,  with  walls  of  medium  thickness,  should  be 
selected.  Experiments  are  often  a  failure  in  thick-walled 
tubes,  because  the  heat  from  the  ordinary  Bunsen  flame  is 
not  able  to  penetrate  sufficiently.  On  the  other  hand,  a  thin- 
walled  tube  is  liable  to  blow  out  when  hot,  if  the  gas  pressure 
inside  is  greater  than  that  of  the  atmosphere  outside. 

Clay  wasting-dishes  are  very  handy  for  muffle  work,  and 
Morgan's  No.  1  is  a  useful  size. 


16—2 


APPENDIX 


TABLE  OF  ELEMENTS,  WITH  SYMBOLS  AND  ATOMIC 
WEIGHTS 

NON-METALS. 


NAME. 

SVMBO,.    Co*'?,  j 

NAME. 

SYMBOL. 

ATOMIC 
WEIGHT. 

Argon    

A 

40'0 

Iodine  

I 

126-54 

Arsenic  ... 

As 

74-9 

Nitrogen 

N 

14-01 

Bromine 

Br 

7976 

Oxygen  

0 

15-96 

Boron    B          109 

Phosphorus      ...  !      P 

30-96 

Carbon  

C         11-97 

Silicon  .. 

Si 

28-0 

Chlorine 

01 

35-37     I  Sclenion 

Se 

78-87 

Fluorine 

F 

19-06    ;|  Sulphur 

S 

31-98 

Hydrogen 

H 

i-o 

Tellurion 

Te 

127-7 

COMMON  METALS. 


Aluminium 

Al 

27-04 

Magnesium 

Mg 

23-94 

Antimony 

Sb 

119-6 

Manganese 

Mn 

54-8 

Barium... 

Ba 

136-86 

Mercury           ...  i     Hg       199'8 

Bismuth 

Bi 

207-5 

Molybdenum    ...  '     Mo         95  -9 

Cadmium 

Cd 

111-7 

Nickel  

Ni 

58-6 

Calcium 

Ca 

39-91 

Platinum 

Pt 

194-3 

Chromium 

Cr 

52-45 

Potassium 

K          39-03 

Cobalt  ... 

Co 

58-6 

Silver    

Ag     i  107-66 

Copper  

Cu 

63-18 

Sodium  

Na         22-995 

Gold      

Au 

196-7 

Strontium 

Sr 

87-3 

Iron 

Fe 

55-88 

Tin        

Sn 

117-35 

Lead     

Pb 

206-92 

Zinc       

Zn         64-88 

REMARKS. — Sb  from  stibium  ;  Hg  from  hydrargyrum  ;  K  from  kalium  ; 
Cu  from  cuprum  ;  Ag  from  argentum  ;  Au  from  aurum  ;  Na  from 
natrium  ;  Fe  from  ferium  ;  Sn  from  stannum  ;  Pb.from  plumbum.  The 
names  of  the  more  important  elements  are  printed  in  thick  type. 

244 


APPENDIX 


245 


RARE  METALS. 


NAME. 

SYMBOL. 

ATOMIC 
WEIGHT. 

NAME. 

SYMBOL. 

ATOMIC 
WEIGHT. 

Beryllium 

Be 

9-08 

Ruthenium 

Ru 

103-5 

Caesium   . 

Cs 

1327 

Samarium 

Sm 

150-0 

Cerium  ... 

Ce 

142-2 

Scandium 

Sc 

43-97 

Didymium 
Erbium  ... 

D 
E 

145-0 
166-0 

Tantalum 
Terbium 

Ta 
Tr 

182-0 
160  0 

Gallium 

Ga 

69-9 

Thallium 

Tl 

203-7 

Germanium 

Ge 

72-32 

Thorium 

Th 

231-96 

Indium  ... 

In 

113-4 

Thulium 

Tu 

169-4 

Iridium  

Ir 

192-5 

Titanium 

Ti 

48 

Lanthanum 

La 

138-5 

Tungsten 

W 

183-6 

Lithium 

Li 

7-01 

Uranium 

U 

239-8 

Niobium 

Nb 

93-7 

Vanadium 

V 

51-1 

Osmium 

Os 

191-0 

Ytterbium 

Yb 

172-6 

Palladium 

Pd 

106-2 

Yttrium 

Y 

89-9 

Rhodium 

Rh 

103-1 

Zirconium 

Zr 

90-4 

Rubidium 

Rb 

85'2 

CORRECTION  OF  THE  VOLUME  OF  A  GAS  FOR  TEMPERATURE, 
PRESSURE,  AND  WATER  VAPOUR. 

Correction  for  Temperature.— Gases  expand  Y|^  of 
their  volume  for  every  increase  of  1°  C.  in  their  temperature 
and  contract  in  the  same  degree  when  their  temperature  is 
decreased  (law  of  Charles).  Now,  suppose  the  gas  to  be  cor- 
rected were  cooled  to  0°  C.,  its  volume  divided  into  273  parts, 
and  the  whole  again  raised  a  degree  at  a  time  to  the  original 
temperature.  Then  for  a  rise  of  1°  the  increased  volume 
would  be  represented  by  274,  for  5°  by  278,  for  t°  by  273 +  /. 
Therefore  the  ratio  between  the  volume  of  a  gas  at  /°  C.  and 
its  volume  at  0°  C.  is  expressed  by  (273 +  £)  :  273,  and  the 
volume  at  0°  C.  is  found  by  multiplying  the  volume  at  t°  C.  by 

273 
•>734-/-     -^  the  temperature  of  the  gas  is  below  0°  C.  t  is 

negative,  and  it  is  to  be  subtracted  from,  and   not   added 
to,  273. 


246  METALLUEGICAL  CHEMISTKY 

Correction  for  Pressure. — The  volume  of  a  gas  is  in- 
versely as  the  pressure  it  sustains,  or  the  volume  decreases  in 
the  same  proportion  as  the  pressure  increases,  and  vice  versa 
(law  of  Boyle  and  Marriott).  The  pressure  to  which  a  gas 
is  subjected  when  standing  over  a  liquid  with  the  liquid  at  the 
same  level  inside  and  outside  the  containing  vessel,  is  exactly 
that  of  the  air  outside,  and  can  be  determined  in  millimetres 
of  mercury  by  reading  the  barometer.  If  the  normal  pressure 
be  taken  as  equal  to  760  millimetres  of  mercury,  and  the 
pressure  of  the  gas  to  be  corrected  is  equal  to  p  millimetres, 

^7  A  r\ 

then  the  ratio  between  the  volumes  at  p  and  at  760  is  — .  and 

p  ' 

the  corrected  volume  is  found  by  multiplying  the  volume  at 
p  by  Tj^T)-    p  may  be  greater  or  less  than  760. 

Correction  for  Water  Vapour.— A  gas  standing  over 
water  becomes  saturated  with  the  vapour  of  the  liquid,  but 
the  weight  of  the  volume  of  water  vapour  present  depends 
upon  the  temperature  of  the  gas.  Thus,  at  a  given  tempera- 
ture a  definite  weight  of  water  gas  will  mix  with  the  other 
gas,  and,  by  Dalton's  law  of  partial  pressures,  will  exert  its 
own  share  of  the  total  pressure  of  the  moist  gas.  But  when 
two  gases  are  mixed  together  their  partial  pressures  are  pro- 
portional to  their  original  volumes.  Therefore,  if  the  partial 
pressure  of  water  vapour  at  the  temperature  of  the  moist  gas 
is  known,  the  volume  of  the  water  vapour  can  be  calculated. 
A  diagram  showing  the  pressure  of  saturated  water  vapour 
for  a  limited  range  of  temperature  is  given  on  next  page. 

The  temperature  of  the  gas  is  to  be  taken  in  degrees  Centi- 
grade, and  the  tension  of  the  vapour  at  that  temperature  is 
expressed  in  millimetres  of  mercury. 

The  temperature  is  measured  on  the  horizontal  lines,  and 
the  tension  on  the  vertical  ones.  The  curve  is  obtained  by 
marking  off  on  the  diagram  the  results  of  a  number  of  experi- 
ments for  determining  the  vapour  tension  at  varying  tern- 


APPENDIX 


247 


peratures,  between  the  two  extreme  temperatures,  and  joining 
the  point  by  a  continuous  line. 

DIAGRAM  SHOWING  THE  TENSION  OF  AQUEOUS  VAPOUR  IN 

A  GAS  SATURATED  WITH  THE  VAPOUR  AT  A 

GIVEN  TEMPERATURE. 


24 


Ul    /J 


5678 


11     12     13    1*     IS    16     17    1&     13    20    21     22   23    24    25 
DEGREES   CENTIGRADE 


To  use  the  diagram,  read  the  temperature  of  the  saturated 
gas  to  the  nearest  tenth ;  find  this  temperature  on  the  hori- 
zontal line ;  estimate  with  the  eye  the  point  on  the  curve 
vertically  above  it,  and  then  the  projection  of  this  point  on 
the  vertical  line  represents  the  required  tension  in  milli- 
metres. Example:  If  the  temperature  is  equal  to  16'3°  C., 
the  tension  is  equal  to  13'8  millimetres  of  mercury. 


INDEX 


PAGE 

Acm,  hydrochloric   .-.          ..  69 

—  nitric          .-.          ..      ~v.  78 

—  sulphuric  . .          . .          . .  63 

Acid-forming  oxides             . .  137 

Acids,  action  of.  on  allovs    . .  106 

—  iron        v.  "  ,  ..  105 

—  basicity  of             ..          ..  153 

—  constitution  of      . .          . .  152 

—  phosphoric             . .          . .  224 

—  silicic         . .          . .          . .  228 

Action,  electro-chemical       . .  107 

Air  gas,  composition  of        . .  180 

Alkalies,  fixed           . .          . .  83 

—  volatile 83 

Allotropic  modification          57,  166 

Alloys,  action  of  acids  on    . .  106 

Aluminium 90 

Amalgamated  zinc     . .          . .  108 

Ammonia        . .          . .          . .  82 

Ammonium  chloride             . .  81 

Analysis          . .          . .          . .  67 

Anhydride 138 

Apatite            222 

Atom,  definition  of  . .          . .  42 

Atomic  heat 128 

—  weights      . .          . .          . .  125 

Atomicity       131 

Atoms  and  molecules,  theory 

of 51 

Avogadro,  law  of       . .          . .  43 


BALANCE 

Basic  oxides 

Basicity  of  acids 

Battery,  voltaic 

Black  oxide  of  manganese 

Bone  ash 


237 
136 
153 
112 
72 
221 


PAGE 

Boyle  and  Marriott,  law  of . .  246 

Brin's  process             . .          .  •  202 

Bromine          . .          . .          . .  76 

Bunsencell 112 

Burette                                   . .  240 


CALORIE 
Carbide  of  iron 
Carbon 

—  dioxide       . .      "•  Y 

—  monoxide  . .        .'. 

—  reduction  by 
Carbonates  and  acids 

—  natural 

—  prepared 


219 
105 
163 
172 
174 
185 
171 
167 
168 


Carbonic  oxide,  reduction  by     190 
Cast  iron         . .          . .          . .     105 

Cell,  Bunsen 112 

Change,  chemical      . .          . .         4 

—  endothermic          . .        177,  214 

—  exothermic  . .        177,  214 

—  physical     ..          . .          ..         3 

Charcoal          ..         ..          ..164 

T—  combustion  of       . .          . .     165 

Charles,  law  of          . .          . .     245 

Chemical  change        . .          . .         4 

—  combination,  laws  of          47,  52 

—  equation    . .          . .      49,  54,  67 

—  equivalents  . .         . .     116 

table  of  . .          . .     125 

Chlorate  of  potassium          . .       75 
Chloride  of  ammonium         . .       81 
Coal ..177 

—  gas 179 

Colloids  228 

Colouring  of  massicot  ..        10 

Combustion    . .          . .          . .     200 


249 


250 


INDEX 


Combustion,  explosive 

PACK 

..     206 

FELSPAR 

PAGE 

..     234 

—  inversion  of 

..     207 

Fixed  alkalies 

..       83 

—  of  charcoal 

..     165 

Flame  

..     208 

Common  salt 

37,69 

—  luminosity  of 

..     210 

Complex  oxides 

..     137 

—  structure  of 

..     209 

—  silicates 

..     229 

Flasks,  conical 

..     243 

Composition,  percentage 

..       65 

Fritting 

..     228 

Compound,  definition  of 

..41,47 

Fusibility  of  silicates 

..     229 

Conical  flasks 

..     243 

Conservation  of  energy 

211,  212 

GAS,  coal 

..      179 

matter 

..       35 

Gases   

3 

Constitution  of  acids 

..     152 

—  diffusion  of 

..       31 

Copper,  oxidation  of 

..       10 

Glass    

..     232 

Crystallization,  water  of 

..     147 

Gold,  subdivision  of  .  . 

..       36 

Crystalloids 

..     228 

Graham 

.  .     227 

Crystals 

..     149 

Graphite 

..     163 

Crucibles,  porcelain  .  . 

..     243 

Current,  electric 

..     109 

HEAT,  action  of,  on  salts 

..     157 

Cyanide,  reduction  by 

..     195 

—  water 

..       30 

—  atomic 

..     128 

DENSITY 

..       33 

—  definition  of 

..     211 

—  of  vapour 

..     126 

—  measurement  of    .  . 

..     214 

Dialysis 

..     228 

—  reduction  by 

.  .     182 

Diffusion  of  gases 

..       31 

—  specific 

..     217 

Dioxide,  carbon 

..     172 

Hydrochloric  acid     .  . 

..       69 

—  sulphur 

94 

and  metals 

.  .  86-91 

Dissociation    ..          ..82 

,  91,  101 

Hydrogen,  preparation  of 

24,94 

Divisibility  of  matter 

36 

—  properties  of 

..       25 

Double  salts   .. 

.  .     155 

—  reduction  by 

..     188 

Dressing 

9 

IGNITION  point 

..     201 

ELECTRIC  current 

..     109 

Indicator,  methyl-orange 

..      141 

action  of,  on  water 

..      27 

—  litmus 

..      136 

Electricity,  reduction  by 

..      184 

Inversion  of  combustion 

..     207 

Electro-chemical  action 

..     107 

Iodine  .  . 

..       76 

Electrolysis 

111 

Ions 

.     Ill 

Electrolyte 

..     Ill 

Iron      

..       24 

Element,  definition  of 

41 

—  action  of  acids  on 

.  .     105 

Elements,  valency  of 

..     131 

—  carbide 

..     105 

Endothermic  change 

177,  214 

—  cast 

..     105 

Energy 

4 

—  oxidation  of 

..       12 

—  conservation  of     .  . 

211,  212 

—  pyrites 

..       61 

—  kinetic 

..     211 

—  rusting  of  .  . 

15 

—  potential 

..     212 

—  wrought 

..     105 

Equation,  chemical  .  .      4 

9,  54,  67 

Isomorphism 

128,  156 

—  thermo-chemical  .  . 

..     201 

Equivalents,  chemical 

..     116 

KINETIC  energy 

..     211 

-  table  of 

..     125 

Exothermic  change  .  . 

177,  214 

LAW  of  Avogadro 

..       43 

Explosive  combustion 

..     206 

—  Boyle  and  Marriott 

.  .     246 

Extraction  of  phosphorus 

..     221 

—  Charles 

..      245 

INDEX 

PAGE 

Laws  of    chemical  combina- 

Oxidation of  lead 

tion 

47,52 

mercury 

Lead,  oxidation  of     .  . 

8 

tin 

-  red  

10 

Oxide,  nitric  .  . 

Levigation 

9 

—  nitrous 

Litmus  indicator 

136 

Oxides,  acid-forming 

Liquids 

2 

—  basic 

Luminosity  of  flame 

210 

—  complex 

—  nomenclature  of   .  . 

MAGNESIUM    

23 

—table  of 

Malleability    .. 

8 

Oxygen 

Manganese,  black  oxide  of  .  . 

72 

—  preparation  of 

Mass 

33 

—  properties  of 

Massicot,  colouring  of 

10 

Matter,  conservation  of 

35 

PEROXIDE,  nitric 

—  definition  of 

33 

Phosphoric  acid 

—  divisibility  of 

36 

Phosphorus     .  . 

—  physical  states  of  .  . 

2 

—  extraction  of 

Measurement  of  heat 

214 

—  red  

temperature 

215 

—  yellow 

Mechanical  mixtures 

50 

Physical  change 

Mercury,  oxidation  of 

17 

—  states  of  matter    .  . 

—  red  oxide  of 

38 

Preparation  of  oxygen 

Metals,  action  of,  on  water  .  . 

22 

Prepared  carbonates  .  . 

—  and  hydrochloric  acid        86,  91 

Porcelain  crucibles     .  . 

nitric  acid 

96 

Potassium  chlorate    .  . 

sulphuric  acid 
—  reduction  by 
Methyl-orange 

91 
193 

83 

Potential  energy 
Properties  of  oxygen 
Pyrites,  iron 

—  orange  indicator 

141 

Metric  System 

236 

REACTION,  reduction  by 

Mixtures,  mechanical 

50 

Red  lead 

Molecule,  definition  of 

42 

—  phosphorus 

Molecules  and  atoms,  theory 

—  oxide  of  mercury  .  . 

of     .  . 

51 

Reduction  by  carbon 

Monoxide,  carbon 

174 

carbonic  oxide 

NATURAL  carbonates 

167 

cyanide 
electricity 

Nitric  acid 

78 

heat 

and  metals 

96 

hydrogen 

—  oxide 

100 

metals  .  . 

—  peroxide 

100 

reaction 

Nitrogen         

20 

Roasting 

Nitrous  oxide 

99 

Rusting  of  iron 

Nomenclature  of  oxides 

138 

salts 

151 

SALT,  common 

silicates 

233 

Saltpetre 

Salts,  action  of  heat  on 

OXIDATION 

19 

—  double 

—  of  copper    .  . 

10 

—  nomenclature  of  .  . 

iron 

12 

—  table  of 

251 


PAGE 
8 

.  17 

.  12 

.  100 

99 

.  137 

.  136 

.  137 

.  138 

.  160 

.  20 

.  202 

.  204 

.  100 

.  224 

.  222 

.  221 

.  222 

.  222 

3 

2 

.  202 

.  168 

.  243 

.  75 

.  212 

.  204 

.  61 

.  196 

.  10 

.  222 

.  38 

.  185 

.  190 

.  195 

.  184 

.  182 

.  188 

.  193 

.  196 

.  61 

.  15 

37,69 

.  78 

.  157 

.  155 

.  151 
160 


252 


INDEX 


Scales,  thermometric 

PAGE 
216 

Thermometric  scales 

PAGE 

2L6 

Silica 

226 

TSu     nvirlfl  tinn   r\£ 

19 

Silicates 

228 

1  Hi,    UAlU.dLJ.UJJl   Ul             •  •                 •  • 

J.Z 

—  complex 

229 

VALENCY  of  elements 

131 

—  fusibility  of 

229 

—  table  of 

133 

—  nomenclature  of   .  . 
-  table  of      

233 
233 

Vapour  density 
water,  correction  for  .  . 

126 
246 

Silicic  acids 

228 

Volatile  alkali           .  . 

83 

Silicon             
Solids  

226 

2 

Voltaic  battery 

112 

Speciiic  heat 
Structure  of  flame 
Sulphides 

217 
209 
61 

WATER           
—  action  of  electric  current 

21 

97 

Sulphur 
—  dioxide 
Sulphuric  acid 
—  and  metals 

55 
94 
63 
91 

on 
—  heat  on        .  . 
metals  on    .  . 
—  of  crystallization 
—  vapour,  correction  for 

At 

30 
22 

147 

246 

TABLE  of  oxides 
salts      
—  silicates 

160 
160 
233 

Weighing 
Weights,  atomic 
Wrought  iron 

239 
125 
105 

valency 

133 

Temperature,  measurement  of 

215 

YELLOW  phosphorus 

222 

Thermo-chemical  equation  .  . 
Thermometer  

201 
215 

ZINC,  amalgamated 

108 

THE   END 


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